GIFT   OF 
Mrs.   G.  N.  Lems 


GENERAL    CHEMISTRY 

PART  I 

PRINCIPLES   AND 
APPLICATIONS 


BY 
LYMAN    C.    NEWELL,    PH.D.    (JOHNS  HOPKINS) 

PROFESSOR   OF   CHEMISTRY  IN   BOSTON   UNIVERSITY 

AUTHOR  OF  "EXPERIMENTAL  CHEMISTRY,"  "DESCRIPTIVE  CHEMISTRY," 
"INORGANIC  CHEMISTRY  FOR  COLLEGES" 


D.   C.   HEATH  &  CO.,   PUBLISHERS 

BOSTON  NEW  YORK         CHICAGO 


COPYRIGHT,     1914 
BY    LYMAN    C.    NEWELL 


154 


31 


PREFACE 

THIS  book  has  been  written  to  meet  the  demand  for  a  simple 
and  practical  treatment  of  the  principles  and  applications  of 
chemistry.  The  author  realizes  that  demands  are  now  being 
made  of  teachers  of  chemistry  which  would  hardly  have  been 
considered  several  years  ago.  Not  only  must  the  student  be 
taught  the  principles  of  chemistry  and  their  applications  in 
daily  life,  but  he  must  be  taught  in  such  a  way  that,  should 
occasion  arise,  he  can  use  chemistry  in  earning  a  living.  The 
plea  so  long  made  that  chemistry  is  practical  and  useful  is  now 
being  tested.  Consequently  the  author  has  kept  one  thought 
in  mind  during  the  preparation  of  this  book,  viz.:  principles 
and  applications  must  go  hand  in  hand.  Principles  are  the 
foundation  upon  which  applications  rest.  To  teach  either  one 
exclusively  is  hazardous,  for  when  separated  one  is  as  barren 
as  the  other  is  superficial.  The  author  has  aimed  to  include 
in  this  book  not  only  the  principles  of  chemistry  universally 
regarded  as  an  essential  part  of  a  well-rounded  course,  but  also 
numerous  practical  applications. 

First,  the  text  includes  a  clear,  simple  treatment  of  the  facts, 
laws,  theories,  and  principles  that  serve  as  the  foundation  of 
chemistry.  This  material  is  written  in  a  style  which  assumes 
that  the  book  is  to  be  read,  studied,  and  used  by  beginners. 
Special  effort  has  been  made  to  present  very  clearly  such  needful 
topics  as  atomic  weights,  equations,  and  valence.  All  of  this 
material  is  made  more  serviceable  by  a  liberal  selection  of 
original  exercises  and  problems  containing  many  novel  and 
attractive  features.  Teachers  will  find  in  these  exercises 
and  problems  abundant  material  to  drive  home  principles 
and  arouse  interest  in  the  vocational  aspects  of  chemistry. 

M623259 


iv  PREFACE 

Second,  the  text  includes  at  strategic  points  illustrated 
descriptions  of  (a)  chemical  processes  —  adequate  descriptions 
which  can  be  used  by  students  in  actual  study  of  the  manu- 
facture of  gas,  acids,  steel,  cement,  lime,  etc.,  (b)  modern  elec- 
trolytic and  electrothermal  processes  for  making  aluminium, 
sodium  hydroxide,  carborundum,  calcium  carbide,  carbon 
disulphide,  etc.,  (c)  recent  inventions,  such  as  the  oxygen 
helmet,  pulmotor,  and  oxyacetylene  blowpipe,  as  well  as  the 
discoveries  centering  around  radium,  and  (d)  the  relation  of 
these  applications  to  industries  and  commerce.  Furthermore, 
in  view  of  the  wide-spread  interest  in  the  chemistry  of  life  and 
the  home,  almost  an  entire  chapter  (Chapter  XVII)  is  devoted 
to  food  and  nutrition,  while  many  cognate  topics  are  scattered 
throughout  the  text. 

The  experiments  prepared  to  accompany  this  book  are 
published  separately  and  are  referred  to  in  the  text  as  Part  II. 
These  experiments  have  been  selected  and  arranged  with 
exceptional  care.  A  novel  feature  is  the  division  of  the  experi- 
ments into  "regular"  and  "supplementary."  The  regular 
experiments  include  those  needed  by  the  average  class,  while 
the  supplementary  set  includes  additional  experiments  of  vary- 
ing length  and  difficulty.  This  liberal  provision  will  permit 
teachers  to  select  experiments  suitable  for  various  needs. 

The  author  is  grateful  to  many  teachers  and  former  students 
for  helpful  suggestions.  He  is  specially  indebted  to  Mr.  Royal 
M.  Frye,  Boston  University,  1911,  for  assistance  in  reading  the 
manuscript  and  proof  and  to  Mr.  Harold  C.  Spencer,  Boston 
University,  1914,  for  making  most  of  the  drawings. 

L.  C.  N. 
BOSTON,  MASS. 
May,  1914. 


CONTENTS 


CHAPTER  PAGE 

I.    CHEMISTRY  —  SUBSTANCES  —  PROPERTIES  —  CHANGES 

—  CLASSES    OF    PROPERTIES    AND    SUBSTANCES  — 
DISTRIBUTION  OF  ELEMENTS i 

II.  OXYGEN n 

III.  HYDROGEN 22 

IV.  SOME  PROPERTIES  OF  GASES 30 

V.  PROPERTIES  OF  WATER 36 

VI.    COMPOSITION  OF  WATER  —  HYDROGEN  DIOXIDE    ...       49 
VII.    LAW  AND  THEORY  —  LAWS  OF  DEFINITE  AND  MULTIPLE 
PROPORTIONS  —  ATOMIC     THEORY  —  ATOMS     AND 

MOLECULES  —  SYMBOLS  AND  FORMULAS 55 

VIII.    CHEMICAL  REACTIONS,  EQUATIONS,  AND  CALCULATIONS     65 
IX.    CHLORINE  —  HYDROCHLORIC  ACID  —  ACIDS,  SALTS,  AND 

BASES 73 

X.    NITROGEN  —  AMMONIA  —  NITRIC  ACID  AND  NITRATES 

—  NITROGEN  OXIDES 86 

XI.    THE  ATMOSPHERE  —  ARGON  —  LIQUID  AIR     104 

XII.  GAY-LUSSAC'S  LAW  OF  GAS  VOLUMES  —  AVOGADRO'S 
HYPOTHESIS  —  MOLECULAR  WEIGHTS  AND  ATOMIC 
WEIGHTS  —  MOLECULAR  FORMULAS  AND  EQUA- 
TIONS   113 

XIII.  VALENCE  —  EQUIVALENT  WEIGHT 126 

XIV.  SOLUTION  —  ACIDS,  BASES,  AND  SALTS 136 

XV.    CARBON  —  OXIDES      AND      CARBONATES  —  HYDROCAR- 
BONS —  CARBIDES  —  CYANOGEN      154 

XVI.    ILLUMINATING  GASES  —  FLAMES 181 

XVII.    OTHER  CARBON  COMPOUNDS  —  FOOD  AND  NUTRITION  .  193 
XVIII.    SULPHUR  —  SULPHIDES  —  SULPHUR  OXIDES,  Acros,  AND 

SALTS 226 

XIX.    BORON  —  BORAX  —  BORIC  Aero 245 


VI 


CONTENTS 


CHAPTER  PAGE 

XX.    SILICON  —  SILICON  DIOXIDE  —  SILICIC  Aero  AND  SILI- 
CATES —  GLASS 248 

XXI.    CLASSIFICATION  OF  THE  ELEMENTS  —  METALS  AND  NON- 
METALS —  PERIODIC  CLASSIFICATION 259 

XXII.    FLUORINE  —  BROMINE  —  IODINE 266 

XXIII.  PHOSPHORUS  —  ARSENIC  —  ANTIMONY  —  BISMUTH  .    .    .  275 

XXIV.  SODIUM  —  POTASSIUM  —  AMMONIUM  COMPOUNDS    .    .    .  288 
XXV.    COPPER  —  SILVER  —  GOLD 303 

XXVI.    CALCIUM  —  STRONTIUM  AND  BARIUM 318 

XXVII.    ALUMINIUM  —  CLAY  AND  CLAY  PRODUCTS 328 

XXVIII.    IRON  —  NICKEL  AND  COBALT 337 

XXIX.    MAGNESIUM  —  ZINC  —  CADMIUM  —  MERCURY 354 

XXX.    TIN  — LEAD 365 

XXXI.    CHROMIUM  —  MANGANESE 374 

XXXII.    PLATINUM 380 

XXXIII.    RADIUM  AND  RADIOACTIVITY 382 

APPENDIX 388 

INDEX      392 


GENERAL   CHEMISTRY 

PRINCIPLES   AND   APPLICATIONS 
CHAPTER  I 

CHEMISTRY  —  SUBSTANCES  —  PROPERTIES  —  CHANGES  — 

CLASSES   OF  PROPERTIES  AND   SUBSTANCES - 

DISTRIBUTION   OF  ELEMENTS 

1.  Chemistry    deals    chiefly    with    substances,    their 
properties,  and  the  changes  they  undergo.     Since  it  is  a 
branch  of  natural  science,  chemistry  also  treats  of  certain 
laws  and  theories  which  help  us  understand  the  relations 
of  different  substances.     Furthermore,  chemistry  includes 
a  description  of  various  industries  and  a  consideration  of 
numerous  processes  connected  with  life  itself. 

2.  Substances. -- The  world  about  us  is  made  up  of 
substances.     Wood,   glass,  paper,   food,   cloth,   soil,   and 
metals  are  familiar  substances;    air,  water,  and  all  other 
gases  and  liquids  are  also  substances.     We  might  define 
substances  as  the  material  of  which  bodies  are  made. 

3.  Properties  of  Substances.  —  Substances  are  recog- 
nized and  distinguished  by  certain  characteristics  known 
as  properties.     A  given  substance  always  has  a  group  of 
properties,  often  some  distinctive  ones  which  enable  us 
to  identify   the    substance.     Thus,   gasoline  is   a  liquid 
which  yields  an  explosive  vapor.     Many  properties  are 
easily  detected   by  mere   observation,   e.g.   color,   taste, 
odor,  hardness,  and  physical   state    (i.e.  whether  solid, 


2  CHEMISTRY 

liquid,  or  gas).  Other  properties  are  readily  found  by 
more  careful  observation  and  experiment,  e.g.  boiling 
point,  melting  point,  solubility,  conductivity  of  heat  and 
electricity,  and  specific  gravity  (i.e.  relative  weight). 
Still  other  properties  are  found  only  by  special  experi- 
ments; to  this  class  belong  many  of  the  properties  ex- 
hibited when  two  or  more  substances  act  chemically 
upon  one  another. 

4.  Changes  in  Substances.  —  We  know  by  experience 
that  substances  are  undergoing  changes.  Thus,  water 
evaporates,  plants  grow,  food  digests,  and  metals  rust. 
Not  only  are  substances  changing  but  they  can  be  changed 
at  will.  For  example,  the  wood,  paper,  and  coal  lie 
together  unchanged  in  the  stove  until  a  lighted  match  is 
applied,  and  then  all  three  substances  take  fire  and  burn, 
that  is  they  change  into  gases  and  ashes. 

We  know  these  changes  in  substances  take  place,  be- 
cause the  original  substance  with  its  distinctive  properties 
disappears  and  one  or  more  substances  with  new  prop- 
erties appear.  Thus,  water  changes  into  steam  or  ice, 
and  black  coal  can  be  changed  into  invisible  gases  and  a 
white  ash. 

Changes  in  substances  are  shown  by  change  in  proper- 
ties. A  physical  change  is  a  change  in  which  the  sub- 
stance remains  the  same,  though  some  of  its  properties 
are  temporarily  changed.  For  example,  copper  trolley 
wire  has  the  properties  characteristic  of  the  metal  cop- 
per; when  the  dynamo  is  in  operation,  however,  the 
copper  is  no  longer  ordinary  copper,  it  is  electrified  copper 
and  remains  so  until  the  dynamo  is  shut  off,  whereupon 
the  copper  then  possesses  only  its  original  properties. 
The  copper  was  only  physically  changed.  That  is,  some 
of  the  distinctive  characteristics  of  copper  were  changed; 


PHYSICAL    AND    CHEMICAL    CHANGES  3 

the  change  was  temporary,  however,  for  as  soon  as  the 
original  conditions  were  restored,  the  copper  lost  its 
electricity,  so  to  speak,  and  became  copper  as  we  ordi- 
narily know  it.  By  this  physical  change,  the  copper  was 
not  fundamentally  or  permanently  changed,  nor  was  it 
transformed  into  another  substance;  it  merely  acquired 
for  a  time  certain  properties  characteristic  of  copper  under 
certain  conditions.  In  physical  changes,  then,  the  change 
is  only  temporary,  and  the  substance  regains  its  familiar 
properties  as  soon  as  the  original  conditions  are  restored. 
A  chemical  change  is  a  change  in  which  one  or  more 
new  substances  are  formed.  For  example,  if  copper 
wire  is  heated  very  hot,  the  copper  disappears  and  is 
replaced  by  a  black,  brittle  solid,  which  does  not  become 
copper  again  when  removed  from  the  flame  and  cooled. 
The  copper  was  chemically  changed.  That  is,  the  distinc- 
tive characteristics  of  the  copper  were  permanently 
changed,  for  as  soon  as  the  original  conditions  were  re- 
stored, the  copper  did  not  reappear  but  in  its  place  was  a 
new  substance  with  properties  quite  different  from  the 
original  copper.  By  this  chemical  change  the  copper  was 
transformed  into  another  substance.  In  chemical  changes 
the  change  is  permanent,  a'nd  the  original  substance  with 
its  characteristic  properties  is  replaced  by  one  or  more  new 
substances  with  characteristic  properties.  Physical  and 
chemical  changes  are  closely  related  and  it  is  not  always 
easy  to  call  certain  changes  entirely  physical  or  exclu- 
sively chemical.  Most  chemical  changes  involve  some 
kind  of  obvious  physical  change,  e.g.  a  change  in  tem- 
perature or  in  physical  state.  The  test  by  which  we 
determine  whether  the  change  in  substances  is  physical 
or  chemical  is  the  formation  of  new  substances;  in.  phy- 
sical changes  new  substances  are  not  formed,  while  in 


4  CHEMISTRY 

chemical  changes  new  substances  are  formed.  (See  Part 
II,  Exps.  2,  4,  5.) 

Examples  of  familiar  physical  changes  are  the  countless 
number  of  changes  in  physical  state  (i.e.  from  solid  to 
liquid  and  to  gas,  and  vice  versa),  the  electrification  of 
glass,  rubber,  and  metals,  and  the  magnetization  of  iron 
in  a  dynamo  or  magnet.  Examples  of  familiar  chemical 
changes  are  the  rusting  of  iron  and  the  tarnishing  of  other 
metals,  the  burning  of  oil  in  a  lamp  and  the  explosion  of 
gasoline  in  an  automobile  engine,  the  digestion  of  food, 
and  the  combustion  of  wood  and  coal. 

5.  Classes  of  Properties. — Physical  properties  are  those 
that  are  characteristic  of  a  substance  as  we  usually  ob- 
serve it  and  also  those  that  can  be  detected  while  a 
substance  is  undergoing  or  has  undergone  a  physical 
change.  Chemical  properties  are  those  that  are  revealed 
when  a  substance  undergoes  a  chemical  change.  Among 
the  important  physical  properties  are  color,  luster,  odor, 
taste,  crystalline  structure,  conductivity  for  heat  and 
for  electricity,  and  weight.  Chemical  properties  appear 
when  substances  are  subjected  to  the  action  of  light  or 
electricity,  and  especially  when  certain  substances  are 
heated  or  are  brought  into  contact  by  pressing,  mixing,  or 
dissolving.  For  example,  several  physical  properties  of 
copper  are  readily  observed  and  others  can  be  detected 
and  measured  by  causing  the  copper  to  undergo  physical 
changes.  Likewise,  its  chemical  properties  can  be  found 
by  performing  certain  experiments.  Thus,  when  warmed 
with  an  acid,  copper  is  changed  into  a  soluble  blue 
substance,  and  when  heated  with  sulphur,  it  is  trans- 
formed into  a  black  solid.  As  a  result  of  these  experi- 
ments, we  say  copper  interacts  readily  with  warm  acid 
and  with  hot  sulphur.  (See  Part  II,  Exps.  1,  3.) 


CLASSES    OF    SUBSTANCES  5 

6.  Classes    of    Substances.  —  We    divide    substances 
into  three  classes  —  mixtures,  compounds,  and  elements. 
These  classes,  especially  compounds  and   elements,  are 
closely  related   and  are  often  conveniently  studied  to- 
gether.    For  example,  the  element  oxygen  will  soon  be 
studied,  and  at  the  same  time  we  shall  see  that  oxygen 
forms  many  compounds  with  other  elements  and  is  also 
one  ingredient  of  the  mixture  of  gases  known  as  air. 

7.  General  Characteristics  and  Relations  of  Mixtures, 
Compounds,  and   Elements.  —  A    mixture    is   composed 
of  two  or  more  substances  in  varying  proportions.    In 
mechanical  mixtures  the  ingredients  may  often  be  dis- 
tinguished  and   quite   readily   separated.     Paint,    many 
kinds  of  food,  gunpowder,  and  muddy  water  are  examples 
of  mechanical  mixtures.     Mechanical  mixtures  are  com- 
mon in  nature,   e.g.   soil,   clay,   and  many  rocks.     The 
substances   making   up   these   mechanical   mixtures   can 
often  be  seen  with  the  eye  or  a  lens,  and  can  be  more  or 
less  easily  separated  by  such  mechanical  operations  as 
grinding,  sifting,  dissolving,  and  filtering.     A  solution  is 
likewise  a  mixture  which  is  composed  of  two  or  more 
substances  in  varying  proportions,  but  the  ingredients  of 
a  solution  cannot  be  distinguished  nor  separated  by  the 
operations  that  are  effective  in  the  case  of  mechanical 
mixtures.     Many  of  the  properties  of  solutions,  like  those 
of  mechanical  mixtures,  depend  upon  the  nature  and  the 
proportion  of  the  ingredients,  but  solutions  have  other 
properties  which  are  very  important.     Solutions  are  some- 
times called  homogeneous  mixtures,  because  all  parts  are 
alike,  and  in  this  respect  they  differ  quite  markedly  from 
mechanical  mixtures. 

A  compound  is  cojnposed  of  two  or  more  substances, 
but   compounds   differ   fundamentally   from   mechanical 


6  CHEMISTRY 

mixtures  and  solutions.  The  ingredients  of  a  compound 
are  special  substances  known  as  elements,  and  in  a  given 
compound  the  elements  are  always  the  same  and  are 
present  in  an  unvarying  ratio  by  weight.  Thus,  the 
compound  commonly  known  as  salt  is  39.34  per  cent  of 
the  element  sodium  and  60.65  Per  cent  of  the  element 
chlorine.  Moreover,  the  elements  that  make  up  a  com- 
pound are  not  mixed  as  in  the  case  of  mechanical  mixtures 
and  solutions  but  are  united  chemically,  i.e.  they  are  not 
lying  side  by  side  nor  merely  intermingled,  but  chemically 
combined;  a  compound,  to  be  separated  into  its  consti- 
tuents, must  be  subjected  to  conditions  which  will  bring 
about  a  chemical  change.  And  finally,  although  com- 
pounds are  homogeneous,  their  properties  differ  from 
those  of  the  elements  that  compose  them.  Thus,  the 
colorless  liquid,  water,  is  a  compound  of  the  gaseous  ele- 
ments hydrogen  and  oxygen.  There  is  a  very  large 
number  of  compounds,  and  all  consist  of  two  or  more 
elements  chemically  combined. 

An  element  consists  of  a  single  substance.  That  is, 
iron  is  nothing  but  iron,  copper  is  nothing  but  copper, 
and  so  with  sulphur,  oxygen,  and  all  the  other  elements. 
Elements  are  fundamental  substances.  They  are  the 
substances  from  which  compounds  are  made.  A  com- 
pound can  be  decomposed  into  its  constituent  elements, 
but  we  cannot  go  further  back  than  an  element.  For 
most  purposes,  it  will  be  correct  to  regard  elements  as 
the  fundamental  substances  from  which  compounds  are 
formed  and  to  which  compounds  may  finally  be  reduced. 

There  are  about  eighty  elements,  and  for  convenience 
each  element  is  designated  by  a  symbol,  which  is  an 
abbreviation  of  its  name.  The  important  elements  and 
their  symbols  are  given  in  the  accompanying  table. 


ELEMENTS 

TABLE  OF  THE  IMPORTANT  ELEMENTS  * 


Element 

Symbol 

Element 

Symbol 

Aluminium  

Al 

Iron  

Fe 

Antimony  

Sb 

Lead 

Pb 

Argon 

A 

Magnesium 

Mff 

Arsenic  

As 

Manganese 

Mn 

Barium          

Ba 

Mercury 

He 

Bismuth 

Bi 

Nickel 

Ni 

Boron     

B 

Nitrogen  . 

N 

Bromine                  

Br 

Oxygen 

o 

Cadmium 

Cd 

Phosphorus 

p 

Calcium  

Ca 

Platinum  

Pt 

Carbon              

c 

Potassium 

K 

Chlorine 

Cl 

Radium 

Ra 

Chromium  

Cr 

Silicon  

Si 

Cobalt                         

Co 

Silver 

A? 

Copper 

Cu 

Sodium 

Na 

Fluorine  

F 

Strontium  

Sr 

Gold                      

Au 

Sulphur  

s 

Hydrogen 

H 

Tin 

Sn 

Iodine 

I 

Zinc 

Zn 

8.  Distribution  of  the  Elements.  —  Our  knowledge  of 
the  abundance  of  the  elements  is  based  on  a  study  of 
the  atmosphere,  the  ocean,  a  shell  of  the  earth's  crust, 
and  the  human  body.  The  atmosphere  contains  about 
20  per  cent  of  oxygen,  79  of  nitrogen,  and  i  of  argon. 
The  elements  in  the  ocean  are  not  free  like  those  in  the 
atmosphere  but  are  combined  in  various  compounds. 
The  abundance  of  the  elements  is  shown  in  the  accom- 
panying tables. 


*  A  complete  table  of  the  elements  may  be  found  on  the  inside  of  the 
back  cover  of  this  book. 


8  CHEMISTRY 

TABLE  OF  THE  APPROXIMATE  COMPOSITION  OF  THE  OCEAN 


Element 

Per  Cent 

Element 

Per  Cent 

Oxygen 

8^.70 

Sulphur 

Hydrogen  

10.67 

Calcium  

QC 

Chlorine 

2.O7 

Bromine 

OO8 

Sodium  

I.I4 

Carbon  

OO2 

Magnesium 

.14- 

Other  Elements 

traces 

In  the  ten  mile  shell  of  the  earth's  crust  the  chief  ele- 
ments are  combined.     Their  proportions  appear  in  the 
following :  - — 
TABLE  OF  THE  APPROXIMATE  COMPOSITION  OF  THE  EARTH'S  CRUST 


Element 

Per  Cent 

Graphic  Proportion 

47.07 

2Q  Ofi 

Aluminium  
Iron 

7.90 

4-  4^ 



Calcium  
Potassium  

3-44 
2.45 



Sodium  
Magnesium 

2-43 

2  4.O 

— 

Remainder  

1.82 

— 

Many  experiments  show  that  the  per  cent  of  the  ele- 
ments in  the  human  body  is  probably  about  as  follows :  — 
TABLE  OF  THE  AVERAGE  COMPOSITION  OF  THE  HUMAN  BODY 


Element 

Per  Cent 

Element 

Per  cent 

Element 

Per  Cent 

Oxygen 

65.00 

Phosphorus 

1.  00 

Magnesium 

0.05 

Carbon 

1  8.  oo 

Potassium 

o-35 

Iron 

0.004 

Hydrogen 

IO.OO 

Sulphur 

0.25 

Iodine 

trace 

Nitrogen 

3.00 

Sodium 

0.15 

Fluorine 

trace 

Calcium 

2.OO 

Chlorine 

0.15 

Silicon 

trace 

EXERCISES  9 

It  is  evident  from  these  tables  that  only  about  a  dozen 
elements  are  abundant.  A  few  elements  in  both  free 
and  combined  states  provide  most  of  the  substances 
studied  in  chemistry. 

EXERCISES 

1.  State    some    characteristic    properties    of    (a)  glass,    (6)  kerosene, 
(c)  water,  (d)  paper,  (e)  air,  (/)  lead. 

2.  Give,  from  your  own  observation,  three  illustrations  of  (a)  physical 
change  and  (b)  chemical  change. 

3.  Select  the  physical  and  the  chemical  changes  from  the  following: 
(a)  Burning  of  wood,  (b)  melting  of  butter,  (c)  freezing  of  an  ice-cream 
mixture,  (d)  weathering  of  granite,  (e)  tarnishing  of  brass  and  other  metals, 
(/)  formation  of  snow,  (g)  decay  of  food,  (h)  seasoning  of  wood,  (i)  forma- 
tion of  dew,  (/)  disappearance  of  fog,  (k)  drying  of  food,  (/)  fading  of  colored 
cloth,  (m)  burning  of  illuminating  gas,  («)  explosion  of  gasoline,  (0)  melt- 
ing of  a  wax  candle,  (p)  burning  of  a  wax  candle. 

4.  Name  (a)  five  substances  you  know  are  elements  and  (b)  five  you 
know  are  compounds. 

5.  Name  (a)  three  familiar  mixtures  and  (b)  three  familiar  solutions. 

6.  How  can  water  be  distinguished  from  gasoline?     Gold  from  brass? 
Glass  from  sand?    Air  from  illuminating  gas?- 

7.  Name  five  elements  with  which  you  were  familiar  before  beginning 
to  study  chemistry.    Name  the  eight  most  abundant  elements  in  the  earth's 
crust  in  their  order. 

8.  Express  the  relative  abundance  of  the  elements  by  a  diagram. 

9.  How  do  elements  and  compounds  differ?     Could  you  prepare  (a)  a 
compound  from  elements,  (6)  elements  from  a  compound,  (c)  compounds 
from  compounds,  (d)  elements  from  elements? 

10.  Define  and  illustrate  (a)  substance,  (b)  physical  change,  (c)  chemical 
change,  (d)  element,  (e)  compound,  (/)  mixture,  (g)  solution,  (h)  symbol. 

11.  What  is  the  derivation  of  the  words  chemistry  and  alchemy?    (Con- 
sult a  Dictionary  or  a  History  of  Chemistry.) 

PROBLEMS 

[The  Metric  System  of  Weights  and  Measures  is  constantly  used  in  Chemistry, 
and  it  should  be  learned  or  reviewed  at  once.    See  Appendix,  §  i.] 

1.  What  is  the  abbreviation  of   gram,  centigram,  liter,  meter,  cubic 
centimeter,  centimeter,  decimeter,  milligram,  millimeter? 

2.  Express  (a)  i  liter  in  cubic  centimeters,  (b)  2  1.  in  cc.,   (c)  i  meter  in 


io  CHEMISTRY 

centimeters,  (d)  250  cm.  in  dm.,  (e)  i  kg.  in  grams,  (/)  250  gm.  in  mg., 
(g)  56.75  1.  in  cc.,  (/*)  1250  cc.  in  1.,  (i)  i  cc.  in  cu.  m. 

3.  Add  2  kg.,  1.5  dg.,  22  eg.,  14  gm.,  and  7  mg.,  and  express  the  sum 
in  (a)  grams,  (b)  decigrams,  (c)  milligrams. 

4.  How  many  cc.  in  (a)  i  liter,  (b)  i  cu.  dm.,  (c)  i  cu.  m.? 

6.   What  is  the  weight  in  grams  of  (a)  i  liter  of  water,  (b)  250  cc.? 

6.  What  is  the  volume  in  cc.  of    (a)  i  kg.  of  water,  (b)  i  1.,  (c)  i  cu. 
dm.,  (d1)  2.2  lb.? 

7.  If  i  m.  of  magnesium  ribbon  weighs  4  dg.,  how  many  mg.  will  5  cm. 
weigh? 

8.  Into  how  many  pieces  5  cm.  long  can  a  glass  tube  i  m.  long  be  cut? 

9.  A  flask  holds  750  cc.      Express  its  capacity  in  (a)  I.,  (b)  cu.  dm., 
(c)  cu.  mm. 

10.  A  bottle  holds  exactly  1250  cc.      How  many  grams  of  water  will 
fill  it?     How  many  kg.?     How  many  pints?      How  many  1.? 

11.  A  pupil  prepared  five  250  cc.  bottles  of  oxygen  gas.     Express  the 
total  volume  in  (a)  cu.  dm.,  (b)  I.,  (c)  cubic  centimeters. 

12.  A  pneumatic  trough  is  80  mm.  deep,  25  cm.  long,  and  i  dm.  wide. 
How  many  1.  of  water  will  it  hold?     How  many  cc.? 

13.  A  cube   measures  1.5  cm.  on  each  edge.      What  is  its  volume  in 
cu.  mm.? 

14.  How  many  cc.  of  mercury  are  needed  to  fill  a  tube  1.2  sq.  mm.  in 
cross  section  and  1.7  m.  long? 

15.  The  square  lead  of  a  pencil  is  2  mm.  wide,  2  mm.  thick,  and  15  cm. 
long.     What  is  its  volume  in  cu.  mm.? 

16.  A  cylindrical  iron  pipe  is  2  m.  long  and  90  mm.  in  internal  diameter. 
How  many  cc.  of  water  will  it  hold? 


CHAPTER   II 

OXYGEN 

9.  Occurrence.  —  Oxygen  is  the  most  abundant  and 
widely  distributed  of  the  elements.     It  occurs  both  as  a 
free  element  and  as  a  constituent  of  many  compounds. 
Mixed  with  nitrogen  and  small  quantities  of  other  gases, 
it  forms  nearly  2 1  per  cent  (by  volume)  of  the  atmosphere. 
Combined  with  hydrogen,  it  constitutes  88. 8 1  per  cent 
(by  weight)   of  water;    combined  with  silicon  and  cer- 
tain metals,  it  makes  up  nearly  half  of  many  common 
minerals  and  rocks  (8).     Compounds  of  oxygen  with  car- 
bon and  hydrogen  form  a  large  part  of  animal  and  vege- 
table matter.     Thus,  the  human  body  contains  about  65 
per  cent  oxygen,  while  vegetable  matter  contains  about 
40  per  cent. 

10.  Preparation.  —  Oxygen  can  be  prepared  from  its 
compounds  or  from  air.     It  was  first  obtained  by  decom- 
posing a  compound  of  oxygen  and  mercury,  now  called 
mercuric  oxide.     This  compound,  when  heated,  decom- 
poses into  oxygen  and  mercury.     If  the  experiment  is 
performed  in  a  test  tube,  the  oxygen  escapes  as  a  gas  and 
the  mercury  condenses  as  a  film  on  the  upper  part  of  the 
tube.     This  experiment  has  historical  interest,  because  it 
was  first  performed  in  1774  by  Priestley,  the  discoverer  of 
oxygen.     (See  Part  II,  Exp.  8  A.) 

The  gas  is  often  prepared  by  decomposing  other  com- 
pounds of  oxygen,  such  as  potassium  chlorate,  lead  dioxide, 


12 


CHEMISTRY 


barium  dioxide,  and  manganese  dioxide.  Thus,  potas- 
sium chlorate  —  a  compound  of  oxygen,  chlorine,  and 
potassium  —  when  heated  to  a  moderately  high  tempera- 
ture yields  all  its  oxygen,  while  a  white  substance  called 
potassium  chloride  remains.  (See  Part  II,  Exp.  8.) 

Oxygen  can  also  be  prepared  from  water.  When  an 
electric  current  is  passed  through  water  which  contains  a 
little  sulphuric  acid,  the  gases  oxygen  and  hydrogen  are 
liberated  (54),  and  when  sodium  peroxide  is  dropped  into 
water,  oxygen  is  liberated.  (See  Part  II,  Exp.  8  D.) 

Oxygen  is  most  conveniently  prepared  in  the  laboratory 
by  heating  a  mixture  of  potassium  chlorate  and  manganese 
dioxide  in  a  glass  or  metal  vessel,  and  collecting  the  liber- 
ated oxygen  in  a  bottle  by  means  of  a  pneumatic  trough 
(Fig.  i).  (See  Part  II,  Exp.  6  I.) 


Fig.  i  —  Apparatus  for  Preparing  Oxygen. 

11.  The  Preparation  of  Oxygen  illustrates  Chemical 
Change.  —  Let  us  consider  mercuric  oxide.  The  chemi- 
cal change  consists  in  the  decomposition  of  the  compound 
mercuric  oxide  into  the  elements  mercury  and  oxygen. 
This  chemical  change  may  be  compactly  expressed  thus :  - 
Mercuric  Oxide  =  Mercury  +  Oxygen 

(Mercury-Oxygen) 


OXYGEN  13 

Such  an  equation  may  be  read:  Mercuric  oxide  equals 
mercury  plus  oxygen.  The  simple  fact  that  mercuric 
oxide  can  be  decomposed  into  mercury  and  oxygen  is  not 
the  only  reason  for  expressing  the  chemical  change  by  an 
equation.  We  can  also  show  by  experiment  that  when  a 
given  quantity  of  mercuric  oxide  is  entirely  decomposed, 
the  weight  of  the  original  mercuric  oxide  equals  the 
sum  of  the  weights  of  the  mercury  and  oxygen  produced. 
This  equality  of  weights  is  very  important  and  will 
be  discussed  later.  At  present  we  shall  use  equations 
merely  to  emphasize  certain  facts  about  chemical  changes, 
chiefly  the  formation  of  new  substances.  Chemical 
changes  like  that  just  described  are  common,  and  the 
term  decomposition  is  applied  to  them.  Decomposition 
may  be  denned  as  a  chemical  change  in  which  a  com- 
pound is  separated  chemically  into  other  substances  which 
are  elements  or  compounds. 

12.  Physical  Properties.  —  Pure  oxygen  gas  has  no 
color,  odor,  or  taste;  certain  impurities  give  a  slight  odor 
and  taste  to  the  gas  prepared  in  the  laboratory.  It  is  not 
very  soluble  in  water.  Oxygen  is  slightly  heavier  than 
air.  One  liter  of  oxygen  weighs  1.429  grams  when  the 
gas  is  at  the  temperature  of  zero  degrees  as  registered  by 
a  centigrade  thermometer  and  also  under  a  pressure  of  760 
millimeters  as  registered  by  a  barometer  (or  briefly  at  o°  C. 
and  760  mm.).  (The  weight  of  a  liter  of  any  gas,  to  be 
of  use  in  chemistry,  must  be  taken  .when  the  volume  of 
the  gas  is  measured  at  the  standard  temperature  (o°  C.) 
and  pressure  (760  mm.)).  (See  Chapter  IV.) 

If  compressed  and  subjected  to  a  very  low  temperature,  oxygen 
becomes  a  pale  blue  liquid,  which  is  slightly  heavier  than  water; 
at  an  extremely  low  temperature  the  liquid  becomes  a  light  blue 
solid. 


14  CHEMISTRY 

13.  Chemical  Properties.  --  The  chief  chemical  prop- 
erty of  oxygen  is  the  ease  with  which  it  combines  or  inter- 
acts with  other  substances.     Oxygen   forms  compounds 
with  most  elements  and  it  interacts  chemically  with  many 
compounds.     This  combining  or  interacting  is  often  made 
conspicuous  by  the  accompanying   light  and  heat.     At 
ordinary  temperatures  oxygen  unites  slowly  with  most 
elements.     Thus,  metals,  such  as  lead,  zinc,  and  copper, 
tarnish  or  rust  slowly,  i.e.  they  combine  slowly  with  the 
oxygen  of  the  air;   with  phosphorus,  however,  the  chemi- 
cal action  is  quite  rapid,  as  may  be  seen  by  the  glow  and 
fumes  when  the  end  of  a  phosphorus-tipped  match  is 
rubbed,  especially  in  the  dark.     The  chemical  activity  of 
oxygen  at  high  temperatures  is  readily  shown  by  putting 
burning  or  glowing  substances  into  it.     The  action  be- 
comes more  energetic,  and  many  substances  burn  rapidly 
and  brilliantly  in  oxygen.     (See  Part  II,  Exp.  6  II.) 

14.  Test  for  Oxygen.  —  The  conspicuous  behavior  of  a  glowing 
stick  or  burning  substance  when  put  into  oxygen  enables  us  to 
distinguish  oxygen   from   other   gases.     This  critical  examination 
which  is  made  to  establish  the  identity  of  oxygen  is  called  testing 
or  making  a  test.     Each  element  has  properties  which  respond   to 
appropriate  tests.    All  compounds  likewise  behave  in  some  decisive 
way  when  subjected  to  tests. 

15.  The    Chief   Chemical   Property   of   Oxygen  illus- 
trates Chemical  Change.  —  In  the  experiments  described 
in  paragraph  13,   one  feature    is   conspicuous,   viz.    the 
disappearance  of  the  original  substances  and  the  form- 
ation   of    new    substances.      The     chemical    change    in 
the  case  of  the  carbon,  sulphur,  iron,  and  magnesium  is 
the  combining  of  oxygen  with  these  elements.    The  oxy- 
gen is  added  chemically  to  each  element  and  the  product 
is  a  compound  of  the  two  elements.     In  the  case  of  sub- 


OXYGEN  15 

stances  like  wood,  which  is  essentially  a  compound  of  car- 
bon, hydrogen,  and  oxygen,  the  chemical  change  is  similar, 
for  the  oxygen  combines  with  one  or  more  of  the  elements 
in  the  compound. 

The  fact  that  the  chemical  change  just  described  is  a 
combining  of  these  elements  with  oxygen  can  be  easily 
verified.  It  has  been  repeatedly  shown  that  oxygen  is 
one  constituent  of  all  the  products  formed  by  burning 
substances  in  that  gas.  Thus,  carbon  forms  an  invisible 
gas  called  carbon  dioxide,  which  is  a  compound  of  car- 
bon and  oxygen.  Similarly,  sulphur,  iron,  and  magne- 
sium form  compounds  of  these  elements  and  oxygen. 

The  chemical  change  illustrating  the  chief  chemical 
property  of  oxygen  is  called  combination.  As  in  the  case 
of  decomposition,  combination  can  be  expressed  by  an 
equation.  Thus:  — 

Carbon  +  Oxide  =  Carbon  Dioxide 

(Carbon-Oxygen) 

Combination  may  be  defined  as  a  chemical  change  in 
which  two  or  more  elements  or  compounds  combine  to 
form  a  compound. 

16.  Oxidation  and  Oxides.  —  The  special  term  oxida- 
tion is  applied  to  those  cases  of  combination  in  which 
oxygen  combines  with  another  element.  Substances 
which  furnish  the  oxygen  are  called  oxidizing  agents. 
Free  oxygen  and  air  are  oxidizing  agents,  though  the 
oxygen  for  oxidation  is  often  provided  by  compounds  of 
oxygen,  especially  those  that  yield  oxygen  readily,  such 
as  potassium  chlorate.  The  compound  formed  by  the 
union  of  oxygen  and  another  element  is  called  an  oxide 
of  that  element.  Thus,  carbon  forms  carbon  dioxide. 
Oxides  of  different  elements  are  distinguished  by  placing 


1 6  CHEMISTRY 

the  name  of  the  element  (or  a  slight  modification  of  it) 
before  the  word  oxide,  e.g.  magnesium  oxide,  lead  oxide, 
nitric  oxide.  Sometimes  di-,  or  a  similar  numerical 
syllable,  is  prefixed  to  the  word  oxide,  e.g.  manganese 
dioxide,  sulphur  trioxide,  phosphorus  pentoxide.  (See 
Part  II,  Exps.  7,  9.) 

17.  Oxidation  and  Combustion.  —  During  oxidation 
heat  is  liberated,  and  if  the  heat  is  intense,  light  is  also 
produced.  If  oxidation  is  slow,  as  in  the  rusting  of  some 
metals,  the  temperature  of  the  oxidizing  substance  may 
not  rise  appreciably,  because  the  heat  escapes  about  as 
fast  as  it  is  liberated.  Sometimes  the  heat  liberated  dur- 
ing slow  oxidation  cannot  escape  readily,  but  accumulates, 
hastens  the  oxidation,  and  finally  the  temperature  rises  to 
such  a  point  that  the  substance  takes  fire.  Thus,  oily 
rags  carelessly  thrown  aside  by  painters,  hay  stored  in  a 
poorly  ventilated  barn,  and  coal  kept  a  long  time  in  the 
warm  hold  of  a  ship  sometimes  take  fire  without  an 
apparent  cause.  Such  fires,  often  unexpected  and  dis- 
astrous, are  said  to  be  due  to  spontaneous  combustion, 
though  they  are  simply  cases  of  slow  oxidation  which 
becomes  accelerated  by  accumulated  heat.  If  oxidation 
is  rapid  or  proceeds  rapidly,  heat 'is  liberated  quickly,  the 
temperature  rises  suddenly,  and  the  oxidizing  substance 
burns,  often  with  dazzling  light.  This  rapid  uniting 
with  oxygen  is  called  combustion.  In  ordinary  language 
combustion  means  fire  or  burning.  Substances  like  paper, 
wood,  coal,  and  oil,  which  burn  readily,  are  called  com- 
bustible; those  like  water,  sand,  stone,  brick,  glass,  and 
plaster,  which  do  not  burn  at  all,  are  called  incombustible. 
As  usually  used  in  chemistry,  the  term  combustion  means 
rapid  oxidation  accompanied  by  heat  and  light.  Oxygen 
is  essential  to  ordinary  combustion,  and  the  gas  is  often 


OXYGEN  17 

called  a  good  supporter  of  combustion.  If  air  is  excluded 
from  a  fire,  the  fire  goes  out.  When  combustible  sub- 
stances burn,  the  carbon  (of  which  they  wholly  or  partly 
consist)  unites  with  the  oxygen  of  the  air,  thereby  form- 
ing the  invisible  gas  carbon  dioxide,  and  the  chemical 
change  is  attended  by  heat  and  light.  Briefly,  a  burning 
substance  is  uniting  rapidly  with  oxygen.  But  since  air 
is  only  about  one  fifth  oxygen  (the  remainder  being 
chiefly  nitrogen,  which  does  not  support  combustion), 
combustion  is  less  rapid  and  hence  less  vigorous  in  air 
than  in  oxygen.  The  temperature  at  which  combustion 
takes  place  varies  between  wide  limits.  Some  substances, 
like  phosphorus  and  gasoline  vapor,  catch  fire  at  a  moder- 
ate temperature,  while  others  do  not  burn  until  heated 
to  extremely  high  temperatures.  Each  substance  has  its 
own  kindling  temperature,  i.e.  the  temperature  to  which 
it  must  be  heated  before  it  will  catch  fire,  though  this 
temperature  depends  somewhat  on  the  form  of  the  sub- 
stance. Application  of  this  fact  is  seen  in  the  use  of  paper 
and  kindling  wood  in  starting  a  fire  in  a  stove. 

The  correct  explanation  of  fire,  burning,  and  combustion  was 
first  made  by  Lavoisier  (1743-1794).  For  many  years  chemists 
had  believed  that  all  combustible  substances  contained  a  principle 
called  phlogiston,  and  that  when  a  substance  burned,  phlogiston 
escaped.  Very  combustible  substances  were  thought  to  contain 
much  phlogiston,  and  incombustible  substances  no  phlogiston. 
Lavoisier,  in  1775,  proved  by  his  own  and  others'  experiments  that 
phlogiston  did  not  exist,  and  that  ordinary  combustion  is  a  process 
of  combination  with  "a  certain  substance  contained  in  the  air." 
Soon  after  he  identified  this  substance  as  oxygen. 

Lavoisier,  in  1778,  named  the  gas  Qxygen(from  the  Greek  oxus, 
acid,  and  gen,  the  root  of  a  verb  meaning  to  produce) ,  because  he 
believed  from  his  experiments  that  oxygen  was  necessary  for  the 
production  of  acids  —  a  view  now  known  to  be  incorrect. 


i8  CHEMISTRY 

18.  Relation  of  Oxygen  to  Life.  —  Free  oxygen  is  essen- 
tial to  all  forms  of  animal  life.  If  an  animal  is  deprived 
of  air,  it  dies.  By  respiration  air  is  drawn  into  the  lungs; 
here  its  oxygen  is  taken  up  by  the  blood,  which  distributes 
it  to  all  parts  of  the  body.  This  oxygen  slowly  oxidizes 
the  tissues  of  the  body  (see  Hemoglobin,  Chapter  XVII). 
By  this  slow  oxidation  waste  products  are  formed  and 
heat  is  supplied  to  the  body.  One  of  these  waste  products 
is  carbon  dioxide  gas,  which  with  other  gases  is  exhaled 
from  the  lungs.  New  tissue  is  built  up  from  the  food 
we  eat.  The  human  body  resembles  a  steam  engine.  In 
each,  the  oxygen  of  the  air  helps  burn  fuel  largely  composed 
of  carbon.  In  the  engine,  the  products  escape  through 
a  chimney  and  the  heat  produced  is  used  to  form  steam 
which  moves  parts  of  the  machine ;  in  the  body,  the  prod- 
ucts escape  through  the  lungs  and  other  organs  and  the 
heat  keeps  the  body  at  the  temperature  at  which 
it  can  best  perform  its  functions.  (Compare  Food 
as  a  Source  of  Energy,  Chapter  XVII.) 

19.  Uses  of  Oxygen.  —  Oxygen  gas  for  indus- 
trial and  scientific  use  is  stored  under  pressure 
in  strong  metal  cylinders  (Fig.  2).  A  mixture  of 
oxygen  and  hydrogen  gas  or  acetylene  gas  if 
burned  in  a  suitable  apparatus  produces  a  very 
hot  flame.  The  oxy-hydrogen  flame  is  used  to 
Fig.  2.  melt  certain  metals  and  to  produce  the  intense 
-  Oxy-  light  of  the  stereopticon,  while  the  oxy-acetylene 

genCyl-  flame  finc[s  application  in  welding  and  in  burning 
mder. 

apart  heavy  steel  structures,  e.g.  girders  of  bridges 
(26,  195).  Oxygen  gas  is  often  administered  to  persons 
who  are  too  ill  or  weak  to  inhale  the  ordinary  volume 
of  air,  while  liquid  oxygen,  on  account  of  its  low  tempera- 
ture and  oxidizing  power,  is  used  in  the  treatment  of  cer- 


OXYGEN  19 

tain  diseases.  Oxygen  is  now  very  generally  used  in  vari- 
ous forms  of  respiratory  apparatus,  e.g.  the  pulmotor  and 
the  oxygen  helmet.  The  pulmotor  is  essentially  a  pump 
by  which  air  rich  in  oxygen  can  be  forced  into  the  lungs 
at  intervals  approximating  the  normal  rate  of  respiration. 
The  pulmotor  is  used  to  resuscitate  persons  who  have 
been  overcome  by  smoke  or  poisonous  gases  (e.g.  illumi- 
nating gas)  or  who  have  been  rendered  unconscious  by  an 
electric  shock.  Fire  departments,  police  officials,  and 
public  health  officers  are  supplied  with  pulmotors  for 
emergencies.  One  form  of  the  oxygen  helmet  is  shown 
in  Fig.  3.  The  apparatus  consists  of  a  leather  helmet 


Fig.  3.  —  Oxygen  Helmet  showing  the  Parts  in  Position  and  the  Course 
of  the  Gases. 

provided  with  a  series  of  tubes  connecting  the  helmet  with 
a  breathing  bag  (A),  a  cylinder  of  compressed  oxygen 
gas  (B),  and  a  regenerating  can  (C)  containing  pieces  of 
potassium  hydroxide  to  absorb  the  water  vapor  and  car- 
bon dioxide  exhaled  from  the  lungs.  The  helmet,  which 
has  a  mica  window,  is  fastened  securely  upon  the  head 
and  neck,  and  the  rest  is  strapped  over  the  shoulders 


20  CHEMISTRY 

like  a  knapsack  —  the  breathing  bag  on  the  chest  (left) 
and  the  oxygen  and  can  on  the  back  (right) .  The  course 
taken  by  the  gases  is  shown  by  arrows.  The  supply  of 
oxygen  needed  by  the  wearer  (as  well  as  the  circulation  of 
the  gases)  can  be  regulated  by  a  valve  on  the  cylinder  (B) ; 
the  nitrogen  originally  inhaled  and  in  the  apparatus  at 
first  is  breathed  over  and  over.  The  cylinder  contains 
oxygen  sufficient  for  about  two  hours.  Some  forms  of 
apparatus  have  a  mouth-breathing  device  instead  of  a 
helmet.  Provided  with  an  oxygen  helmet  (or  similar  de- 
vice) a  man  can  safely  enter  places  where  the  air  con- 
tains smoke  or  poisonous  gas,  and  make  repairs,  extin- 
guish fires,  or  rescue  workmen  who  Jiave  been  overcome. 
Extensive  use  is  made  of  this  protective  device  in  mine 
disasters,  largely  through  the  efforts  of  the  United  States 
Bureau  of  Mines. 

20.  Ozone  is  a  gas  related  to  oxygen,  though  its  properties 
differ.  It  is  formed  when  electric  sparks  pass  through  the  air,  and 
is  therefore  produced  when  electrical  machines  are  in  operation  and 
during  thunder  storms.  Slow  oxidation,  especially  of  moist  phos- 
phorus, produces  ozone. 

EXERCISES 

1.  Name  several  compounds  from  which  oxygen  can  be  prepared. 

2.  Summarize  the  physical  properties    of    oxygen.     What  is  its  most 
characteristic  chemical  property? 

3.  If  air  contains  a  large  proportion  of  another  gas  besides  oxygen,  how 
must  the  general  properties  of  this  other  ingredient  compare  with  those  of 
oxygen? 

4.  Define  and  illustrate    (a)  oxidation,   (b}  oxide,    (c)  combustion,   (d) 
oxidizing  agent. 

6.  Give  the  name  and  symbol  of  each  element  mentioned  in  studying 
oxygen;  also  the  name  of  each  compound. 

6.  What  general  chemical  change  is  involved  in  ordinary  burning? 
What  class  of  chemical  changes  is  illustrated  by  (a)  preparation  of  oxygen 
from  mercuric  oxide,  (b)  burning  of  sulphur  in  oxygen? 


OXYGEN  21 

7.  Cite    cases    of    so-called    spontaneous   combustion  of    which  you 
have  heard  or  read.     Suggest  methods  to  prevent  spontaneous  combus- 
tion of  (a)  oily  rags,  (b)  coal,  (c)  hay. 

8.  Suggest  an  experiment  to  show  that  air  contains  oxygen. 

9.  In  which  will  a  glowing  piece  of  charcoal  burn  more  vigorously, 
gaseous  or  liquid  oxygen?     Why? 

10.  What  chemical  part  does  oxygen  take  in  (a)  respiration,  (b)  burn- 
ing, (c)  combustion,  (d)  oxidation,  (e)  kindling  a  fire? 

11.  Essay  topics:    (a)  Uses  of  oxygen,     (b)  Discovery  of  oxygen,      (c) 
Priestley,     (d)  Oxygen  and  life,     (e)  Combustion.     (/)  Lavoisier. 

PROBLEMS 

1.  (a)  What  is  the  weight  in  gm.  of  35  1.  of  oxygen  gas?     (b)  Of  35,000 
cc.?     (c)  Of  35  cubic  decimeters? 

2.  How  many  grams  does  a  cubic  meter  of  oxygen  gas  weigh? 

3.  How  many  gm.  of  oxygen  gas  are  there  in  a  bottle  holding  2.5  1.  (at 
o°  C.  and  760  mm.)? 

4.  A  pupil  prepared  enough  oxygen  gas  to  fill  a  tank  holding  i  cu.  m. 
at  o°  C  (and  760  mm.).     How  many  gm.  were  prepared? 

6.    (a)  How  many  liters  (at  o°  C,  and  760  mm.)  will  25  gm.  of  oxygen 
gas  occupy?     (b)  How  many  gm.  will  25  1.  of  oxygen  weigh? 

6.  How  many  kg.  of  oxygen  gas  (at  o°  C.  and  760  mm.)  are  needed  to 
fill  a  tank  measuring  250  m.  X  550  m.  X  1055  m.? 

7.  How  many  gm.  of  oxygen  gas  (at  o°  C.  and  760  mm.)  in  a  cylindrical 
gas  holder  which  is  i  m.  high  and  30  cm.  in  diameter? 

8.  If  air  contains  23  per  cent  of  oxygen  by  weight,  how  many  liters  of 
oxygen  gas  (at  o°  C.  and  760  mm.)  can  be  obtained  from  800  gm.  of  air? 

9.  If  air  contains  21  per  cent  of  oxygen  by  volume,  how  many  gm.  of 
oxygen  gas  can  be  extracted  from  950  1.  of  air? 

10.  A  pupil  prepared  five  bottles  of  oxygen  each  holding  250  cc.  (at 
o°  C.  and  760  mm.).     How  many  gm.  of  oxygen  were  prepared? 

11.  Water  contains  88.82  per  cent  of  oxygen.     Suppose  .5  kg.  was  de- 
composed, how  many  liters  of  oxygen  (at  o°  C.  and  760  mm.)  were  formed? 

12.  Water  contains  88.82  per  cent  of  oxygen.     Suppose  200  liters  of 
oxygen  (at  o°  C.  and  760  mm.)  have  been  obtained;    how  many  gm.  of 
water  were  decomposed? 

13.  A  room  is  10  m.  long,  5  m.  wide,  and  4  m.  high.     How  many  gm. 
of  water,  containing  88.82  per  cent  of  oxygen,  must  be  decomposed  to 
furnish  enough  oxygen  (at  o°  C.  and  760  mm.)  to  fill  the  room? 

14.  Potassium  chlorate  contains  39.18  per  cent  of  oxygen.     If  35  gm. 
are  heated,  how  many  bottles  each  containing  250  cc.  can  be  filled  with  the 
liberated  oxygen? 


CHAPTER   III 


HYDROGEN 

21.  Preparation.  —  Hydrogen  is  readily  prepared  from 
acids.  This  is  accomplished  by  allowing  certain  metals 
and  acids  to  interact.  The  metals 
usually  employed  are  zinc,  iron,  or 
magnesium,  and  the  acids  are  dilute 
water  solutions  of  sulphuric  acid  or 
hydrochloric  acid .  The  hydrogen  comes 
from  the  acid,  and  the  metal  combines 
with  the  rest  of  the  acid  to  form  a 
compound  which  usually  remains  dis- 
solved in  the  apparatus.  In  the  lab- 
oratory hydrogen  is  usually  prepared 
in  a  small  generator,  and  collected 
over  water  in  a  pneumatic  trough. 
On  a  large  scale  a  Kipp  apparatus 
(Fig.  4)  is  sometimes  used.  No  flame 
should  be  near  during  the  preparation, 
because  mixtures  of  air  and  hydrogen 
Fig.  4.— Kipp  Appara-  explode  violently  when  ignited.  (See 

tus    for    Generating     part  jj    £  1Q    12  A  and  R) 

Hydrogen. 

Hydrogen    can    be    obtained    from 

water  by  allowing  certain  metals  and  water  to  interact. 
If  a  small  piece  of  sodium  is  dropped  upon  cold  water, 
the  sodium  melts  into  a .  shining  globule,  which  spins 
about  rapidly  on  the  water  with  a  hissing  sound,  and 
finally  disappears  with  a  slight  explosion.  Calcium  inter- 


HYDROGEN 


acts  slowly  with  water,  but  potassium  interacts  so  rapidly 
that  the  heat  ignites  the  liberated  hydrogen  (Fig.  5) .  (See 
Part  II,  Exp.  12  D  and  E.) 

Hydrogen,  together  with  oxygen, 
is  liberated  from  water  by  passing 
a  current  of  electricity  through 
water  containing  a  little  sulphuric 
acid  (54  i).  Hydrogen  can  also  be 
prepared  by  passing  steam  —  the 
gaseous  form  of  water — over  heated 
metals  (Fig.  6).  This  experiment 
was  first  performed  by  Lavoisier,  in 
1783,  while  he  was  studying  the  composition  of  water. 
He  passed  steam  through  a  red-hot  gun  barrel  containing 


o 


Fig.    5.  —  Interaction    of 
Water  and  Potassium. 


Fig.  6.  —  Modern  Form  of  Lavoisier's  Apparatus  for  Showing  the  Forma- 
tion of  Hydrogen  by  the  Interaction  of  Steam  and  Heated  Iron. 

bits  of  iron.  The  oxygen  of  the  steam  combined  with 
the  iron,  and  the  hydrogen  escaped  from  the  tube.  Since 
Lavoisier  was  then  studying  the  composition  of  water 


24  CHEMISTRY 

and  not  especially  the  properties  of  hydrogen,  he  natur- 
ally thought  of  this  gas  as  essential  for  forming  water. 
So  he  gave  the  gas  the  name  hydrogen,  which  means 
literally  "water  former." 

Hydrogen  can  also  be  prepared  by  boiling  solutions  of 
certain  bases  with  metals,  e.g.  sodium  hydroxide  with 
aluminium.  (See  Part  II,  Exp.  12  C.) 

22.  Chemical  Changes  illustrated  by  the  Preparation 
of   Hydrogen.  —  The   preparation   of   hydrogen   by   the 
interaction  of  a  metal  and  an  acid  illustrates  a  third 
kind  of  chemical  change,  viz.   substitution,  or,  as  it  is 
sometimes  called,  displacement  or  replacement.     In  the 
case  of  zinc  and  sulphuric  acid,  the  hydrogen  is  displaced 
from  the  acid  and  the  zinc  takes  its  place;  i.e.  zinc  is  sub- 
stituted chemically  for  hydrogen.     This  chemical  change 
can  be  expressed  by  the  following  equation:  — 

Zinc  +   Sulphuric  Acid    =  Hydrogen  -f-  Zinc  Sulphate 

(Hydrogen-Sulphur-Oxygen)  (Zinc-Sulphur-Oxygen) 

We  might  define  substitution  as  a  chemical  change  in 
which  one  element  replaces  another  in  a  compound. 

23.  Physical  Properties.  —  Hydrogen  has  no  taste  or 
color.     The  pure  gas  has  no  odor,  though  hydrogen  as 
ordinarily  prepared  has  a  disagreeable  odor,  due  mainly 
to  impurities  in  the  metals  used.     Hydrogen  is  the  lightest 
known  substance.     One  liter  of  hydrogen  at  o°  C.  and 
760   mm.  weighs   only   .0898  gm.      Volume  for  volume 
hydrogen  is  about  one  fourteenth  as  heavy  as  air  and  one 
sixteenth  as  oxygen.     Pure  hydrogen  is  not  poisonous. 
Hydrogen  is  not  very  soluble  in  water,  but  it  is  absorbed 
by  several  metals,  especially  palladium.     The  absorption 
of  gases  by  metals  is  called  occlusion.     Hydrogen  diffuses 
readily;  i.e.  it  quickly  passes  through  porous  substances, 


HYDROGEN 


mixes  rapidly  with  other  gases,  and  freely  escapes  into 
space  in  all  directions.  Hydrogen  has  been  liquefied 
and  solidified.  Both  liquid  and  solid  are  colorless  and 
transparent. 

24.  Chemical  Properties.  —  The  chief  chemical  prop- 
erty of  hydrogen 
is  the  readiness 
with  which  it 
unites  with  some 
elements,  especi- 
ally oxygen  and 
chlorine,  whether 
these  elements  are 
free  or  constitu- 
ents  of  co im- 
pounds. Hydro- 
gen burns  in  air 
and  in  oxygen 
with  an  almost 
invisible  but  very 


Fig.  7.  — Apparatus  for  Burning  Hydrogen. 


hot  flame.  A  platinum  or  copper  wire  held  in  the  flame 
quickly  becomes  red-hot.  If  a  small,  dry,  cold  bottle  is 
held  over  the  flame,  moisture  is  deposited  inside  the 
bottle.  Water  is  the  product  of  the  combustion  of 
hydrogen. 

The  burning  of  hydrogen  and  the  properties  of  the  hydrogen 
flame  can  be  shown  readily  by  experiment.  Hydrogen  is  gener- 
ated, as  usual,  and  passed  through  a  U-tube  containing  calcium 
chloride  to  remove  the  water  vapor  (Fig.  7) .  After  all  the  air  has 
been  driven  out  of  the  apparatus  by  the  hydrogen,  the  gas  is  lighted 
at  the  exit.  Unusual  precautions  must  be  used  to  avoid  an  ex- 
plosion, and  before  the  experiment  is  performed  the  directions  should 
be  looked  up  (see  the  author's  Descriptive  Chemistry,  page  472, 
and  Experimental  Chemistry,  page  340). 


26  CHEMISTRY 

The  film  of  water  that  may  be  seen  on  the  bottom  of  a 
vessel  placed  over  a  lighted  gas  range  or  a  Bunsen  burner 
is  the  condensed  vapor  formed  by  the  burning  hydrogen 
and  hydrogen  compounds  of  the  illuminating  gas.  -Or- 
ganic substances  containing  hydrogen,  such  as  wood  and 
paper,  when  burned,  yield  water  as  one  of  their  products. 

The  chemical  change  that  occurs  in  the  burning  of 
hydrogen  can  be  expressed  thus:— 

Hydrogen  +  Oxygen  =  Water 

(Hydrogen-Oxygen) 

It  is  an  example  of  combination  and  also  of  oxidation; 
the  two  elements  combine  to  form  a  compound,  and 
moreover  one  of  the  elements  (the  hydrogen)  undergoes 
oxidation. 

Although  a  small  jet  of  hydrogen  burns  quietly  in  air 
or  oxygen,  a  mixture  of  hydrogen  and  air  burns  so  rapidly 
that  the  combustion  is  practically  an  explosion.  There- 
fore, the  air  should  be  fully  expelled  from  the  apparatus 
in  which  hydrogen  is  being  generated  and  all  leaky  joints 
should  be  tightened  before  the  gas  is  collected;  no  flames, 
large  or  small,  should  be  near.  Neglect  of  these  pre- 
cautions has  caused  serious  accidents. 

Hydrogen  not  only  combines  energetically  with  free 
oxygen,  but  it  also  withdraws  oxygen  from  compounds. 
This  chemical  removal  of  oxygen  is  called  reduction,  and 
the  substances  that  remove  the  oxygen  are  called  reduc- 
ing agents.  Hydrogen  is  a  vigorous  reducing  agent,  just 
as  oxygen  is  an  energetic  oxidizing  agent.  When  oxides 
of  certain  metals  are  heated  in  a  current  of  hydrogen,  the 
oxygen  of  the  oxide  is  chemically  removed  and  combines 
with  the  hydrogen  to  form  water ;  the  metal  is  left  uncom- 
bined.  Thus,  by  heating  copper  oxide  in  hydrogen, 


HYDROGEN  27 

water  and  metallic  copper  are  produced.  Chemically 
speaking,  the  copper  oxide  is  reduced  by  the  hydrogen. 
The  chemical  change  is  substitution  (the  hydrogen  being 
substituted  chemically  for  the  metal),  and  it  can  be 
expressed  thus:  — 

Hydrogen  +  Copper  Oxide       =        Water  +  Copper 

(Copper-Oxygen)  (Hydrogen-Oxygen) 

This  chemical  change  can  also  be  interpreted  from  the 
standpoint  of  oxidation,  because  the  hydrogen  is  oxi- 
dized to  water  at  the  same  time  the  copper  oxide  is  re- 
duced. In  fact,  the  processes  of  reduction  and  oxidation 
are  closely  related  and  either  one  may  be  emphasized  in 
interpreting  the  chemical  change.  It  is  preferable,  how- 
ever, at  this  stage  to  define  reduction  as  the  removal  of 
oxygen  from  a  compound,  postponing  the  details  of  the 
process  until  more  facts  are  available.  In  its  simplest 
form,  reduction  is  the  opposite  of  oxidation.  (See  Part 
II,  Exp.  11.) 

25.  Test  for  Hydrogen.  —  The   test  for  hydrogen  is 
that  it  extinguishes  a  small  flame,  such  as  a  blazing  taper 
or  joss  stick,  but  is  lighted  at  the  same  time,  often  with 
an  explosion,   and   continues   to  burn  until   the  gas  is 
exhausted. 

26.  Uses  of  Hydrogen.  —  On  account  of  its  extreme 
lightness,  hydrogen  is  used   to   fill 

balloons  and  air  ships.  The  in- 
tense heat  of  the  hydrogen  flame  is 
utilized  in  the  oxy-hydrogen  blow- 

".    i  r     ,v       Fig.    8.  —  Oxy-hydrogen 

pipe.     The    essential    part    of    the         Blowpipe  Burner. 

burner    (Fig.    8)     consists    of    two 

pointed  metal  tubes.     The  inner  and  smaller  one  (A)  is 

for  the  oxygen,  and  the  outer  and  larger  one  (B)  for  the 


28  CHEMISTRY 

hydrogen;  the  gases  are  forced  out  of  these  small  openings 
by  the  pressure  maintained  in  the  storage  tanks.  The 
flame  is  used  to  melt  platinum.  When  the  flame  strikes 
against  a  piece  of  lime,  the  latter  becomes  intensely 
bright.  Thus  used,  it  is  called  the  lime  or  calcium  light, 
and  is  often  employed  in  operating  the  stereopticon. 

27.  Occurrence  of  Hydrogen.  —  Combined  with  oxy- 
gen, it  forms  water  —  its  most  abundant  compound. 
With  carbon,  it  forms  hydrocarbons,  which  are  ingredi- 
ents of  natural  gas,  illuminating  gas,  and  the  products 
from  petroleum  (e.g.  kerosene,  gasoline,  and  paraffin). 
With  carbon  and  oxygen,  it  forms  a  large  number  of  com- 
pounds vitally  connected  with  plants  and  animals,  such 
as  starch,  sugar,  and  fat.  With  nitrogen  it  forms  the 
familiar  compound,  ammonia;  and  with  sulphur,  the  bad 
smelling  gas,  hydrogen  sulphide.  Hydrogen  is  also  an 
essential  constituent  of  all  acids  and  bases. 

EXERCISES 

1.  Name  several  substances  which  contain  hydrogen.  Suggest  a  method 
of  obtaining  hydrogen  gas  from  them. 

2.  Compare  oxygen  and  hydrogen. 

3.  How  can  hydrogen  be  distinguished  from  oxygen? 

4.  Summarize  the  physical  properties  Of  hydrogen.     What  is  its  char- 
acteristic chemical  property? 

5.  Why  is  there  danger  of  an  explosion  in  generating  hydrogen?     How 
can  the  danger  be  avoided? 

6.  Define  and  illustrate  (a)  reduction  and  (b)  reducing  agent.     Are  the 
terms  deoxidize  and  reduce  synonymous? 

7.  Describe  the  oxy-hydrogen  blowpipe.     State  the  use. 

8.  In  using  hydrogen  for  balloons,  what  property  of  the  gas  might  cause 
disaster? 

9.  What  chemical  change  occurs  when  hydrogen  burns  in  air? 


HYDROGEN  29 


PROBLEMS 

1.  Sulphuric  acid  contains  2.04  per  cent  of  hydrogen,     (a)  How  many 
gm.  must  be  decomposed  to  yield  85  liters  of  hydrogen  (at  o°  C.  and  760 
mm.)?     (b)  85  cc.?     (c)  How  many  liters  of  hydrogen  gas  (at  o°  C.  and 
760  mm.)  can  be  prepared  from  750  gm.  of  sulphuric  acid?     (d)  How  many 
cc.  of  gas? 

2.  Hydrochloric  acid  contains  2.74  per  cent  of  hydrogen.     How  many 
gm.  must  be  decomposed  to  yield  (a)  44  liters  of  hydrogen  (at  o°  C.  and 
760  mm.)?     (b)  44  cc.?     (c)  How  many  cc.  of  gas  (at  o°  C.  and  760  mm.) 
can  be  made  from  70  gm.  of  the  acid?     (d)  How  many  liters  of  gas? 

3.  Water  contains   11.18  per  cent  of  hydrogen.     How  many  cubic 
centimeters  of  hydrogen  gas  (at  o°  C.  and  760  mm.)  can  be  prepared  from 
230  gm.?     How  many  liters? 

4.  (a)  How  many  liters  (at  o°  C.  and  760  mm.)  will  45  gm.  of  hydrogen 
gas  occupy?     (b)  How  many  grams  will  45  liters  weigh? 

5.  A  vessel  is  i  m.  long,  20  cm.  wide,  and  2.5  dm.  high.     How  many 
grams  of  hydrogen  gas  (at  o°  C,  and  760  mm.)  will  it  contain? 

6.  A  room  is  10  m.  long,  10  m.  wide,  and  7  m.  high,     (a)  How  many 
gm.  of  hydrogen  (at  o°  C.  and  760  mm.)  will  fill  it?     (b)  How  many  kilo- 
grams? 

7.  A  student  prepared  enough  hydrogen  gas  (at  o°  C.  and  760  mm.)  to 
fill  six  bottles,  each  holding  250  cc.     How  many  gm.  were  prepared? 

8.  A  cylindrical  tank  i  m.  long  and  25  cm.  in  diameter  is  filled  with 
hydrogen  (at  o°  C.  and  760  mm.).     How  many  gm.  does  the  gas  weigh? 

9.  A  German  hydrogen  manufactory  produces  daily  19,000  cubic  meters 
of  gas.     (a)  Express  this  volume  in  liters  and  in  cc.     (b)  What  would  this 
volume  weigh,  if  measured  at  o°  C.  and  760  mm.? 

10.  One  of  the  Zeppelin  balloons  (recently  destroyed)  had  a  capacity 
of  351,150  cubic  feet.  What  weight  of  hydrogen  was  needed  to  fill  it? 
(Assume  i  cu.  m.  =  35.32  cu.  ft.;  also  that  the  volume  of  hydrogen  was 
measured  at  o°  C.  and  760  mm.) 


CHAPTER   IV 

SOME   PROPERTIES    OF  GASES  — WEIGHT  OF  A  LITER 
OF  OXYGEN 

28.  Introduction.  —  Gases  behave  quite  uniformly  with 
changes   in  temperature  and  pressure.      In  this   chapter 
we  shall  study  the  effect  of  changes  of  temperature  and 
pressure  upon  the  volume  of  a  gas. 

29.  Relation   of    Gas   Volumes   to    Temperature    and 
Pressure.  --  The  actual  volume  occupied  by  a  gas  de- 
pends upon  the  temperature  and  pressure  prevailing  at 
the   time   of  observation.     A  gas  expands  with  rise   of 
temperature  or  with  decrease  of  pressure;    it  contracts 
with  fall  of  temperature  or  with  increase  of  pressure. 
The  normal  or  standard  temperature  at  which  gases  are 
measured  is  zero  degrees  on  the  centigrade  thermometer 
(or  briefly  o°  C.),  and  the  normal  or  standard  pressure  is 
the  pressure  indicated  by  the  barometer  when  the  mer- 
cury column  is  760  millimeters  high  (or  briefly  760  mm.). 
Under  these  conditions,  which  are  called  standard  con- 
ditions, a  liter  of  oxygen  gas  weighs  1.429  gm.     But  at 
another  temperature  or  pressure  the  liter  would  contain 
a  different  quantity  of  oxygen  gas,  and  would  therefore 
have  a  different  weight.     For  example,  if  the  pressure  is 
increased,  the  volume  becomes  less,  more  gas  must  be 
added  to  bring  the  volume  up  to  a  liter,  and  this  second 
liter  of  oxygen  would  weigh  more  than  1.429  gm.     That 
is,  a  liter  vessel,  when  full,  always  contains  a  liter,  but 
the  weight  of  the  contents  varies  with  the  quantity  of 


SOME  PROPERTIES  OF  GASES  31 

gas  contained  in  this  volume.  Clearly,  if  we  wish  to  find 
the  weight  of  a  given  volume  of  a  gas  or  to  compare  the 
weights  of  gases  by  means  of  their  volumes,  as  we  often 
do  in  chemistry,  we  must  know  the  conditions  under  which 
the  volume  is  measured.  If  all  gases  could  be  measured 
at  o°  C.  and  760  mm.,  their  volumes  would  be  comparable, 
and  the  weights  deduced  or  obtained  directly  from  these 
volumes  would  be  a  true  measure  of  the  actual  quantity 
of  the  gases  in  the  observed  volumes.  But  it  is  incon- 
venient to  measure  gases  experimentally  at  o°  C.  and  760 
mm.  So  it  is  customary  to  measure  the  volume  under 
the  conditions  existing  at  the  time  of  the  experiment, 
and  then  reduce  the  observed  volume  to  the  volume  it 
would  occupy  under  standard  conditions. 

30.  Relation  of  Gas  Volumes  to  Changes  in  Temperature. 
-If  a  gas  at  o°  C.  is  heated  to  i°  C.,  it  expands  1/273 
of  its  volume;  heated  to  5°  C.,  or  to  25°  C.,  or  to  100°  C., 
it  expands  5/273,  or  25/273,  or  100/273,  and  so  on. 
Similarly,  if  it  is  cooled  from  o°  C.  to  -  i°  C.,  it  contracts 
1/273,  and  so  on.  We  can  state  these  relations  between 
temperature  and  volume  in  another  way.  If  we  let  273 
represent  the  volume  of  the  gas  at  o°  C.,  then  its  volumes 
at  the  temperatures  given  above  would  be  represented 
by  274,  278,  298,  373,  272  respectively,  and  similarly 
for  other  temperatures.  This  regularity  of  behavior  is 
shown  by  all  gases,  and  may  be  stated  thus:  All  gases 
under  constant  pressure  expand  or  contract  equally  for  equal 
change  of  temperature.  General  statements  like  this, 
which  summarize  related  facts,  are  called  laws.  This  law 
is  known  as  the  law  of  Charles  from  the  man  (Charles, 
1746-1822)  who  proposed  it. 

The  application  of  the  law  of  Charles  to  the  reduction  of  a  gas 
volume  to  the  volume  it  would  occupy  at  o    C.  can  be  readily  under- 


32  CHEMISTRY 

stood  by  an  example.  Suppose  10  liters  of  oxygen  gas  at  15°  C. 
are  to  be  reduced  to  the  volume  it  would  occupy  at  o°  C.  Let  the 
volume  at  o°  C.  be  represented  by  273;  then  the  volume  at  15°  C. 
would  be  represented  by  273  +  15.  But  273  +  15  and  273  are  in 
the  same  ratio  as  10  (the  known  volume  at  15°  C.)  and  X  (the  un- 
known volume  at  o°  C.).  Therefore  we  can  state  these  relations 
in  a  proportion,  thus:  — 

273  +  15  :  273  :  :  10  :  x;    x  =  9.479  liters. 

Therefore  10  liters  of  oxygen  at  15°  C.  would  occupy  9.479  liters  at 
o°  C.  The  mathematical  operation  of  finding  the  volume  a  gas 
would  occupy  at  o  C.  is  called  reducing  to  standard  temperature  or 
correcting  for  temperature.  Since  t  can  be  substituted  for  any 
temperature  (above  or  below  o°  C.),  the  general  form  of  the  propor- 
tion can  be  written:  — 

2  73  +  t  :  273  ::  known  vol.  :  vol.  at  o°  C. 

31.   Relation  of  Gas  Volumes  to  Changes  in  Pressure. 

—  It  is  found  by  experiment  that  the  volume  of  a  gas  at  a 
constant  temperature  is  inversely  proportional  to  the  pressure. 
This  statement  is  the  law  of  Boyle,  because  it  was  first 
announced  by  Boyle  (1626-1691)  about  1660.  Boyle's 
law  means  that  the  greater  the  pressure,  the  less  the 
volume,  and  vice  versa.  The  normal  pressure,  as  stated 
above,  is  760  mm.  Gases,  as  a-  rule,  are  collected  and 
confined  over  water  or  some  other  liquid  whose  surface 
is  exposed  to  the  atmosphere,  and  since  atmospheric 
pressure  is  transmitted  through  the  liquid  to  the  gas,  the 
pressure  which  the  gas  is  under  is  found  by  reading  the 
pressure  recorded  by  the  barometer  at  the  time  the  gas 
volume  is  read.  The  reduction  of  the  observed  volume 
to  the  volume  it  would  occupy  at  760  mm.  is  performed 
by  applying  Boyle's  law. 

An  illustration  will  make  the  application  clear.    Suppose  we  have 
10  liters  of  oxygen  gas  at  775  mm.  and  wish  to  know  its  volume  at 


SOME  PROPERTIES  OF  GASES  33 

760  mm.  According  to  Boyle's  law,  gas  volumes  are  inversely 
proportional  to  the  pressures;  i.e.  the  observed  pressure  bears  the 
same  relation  to  the  normal  pressure  as  the  normal  volume  bears  to 
the  observed  volume.  Applying  these  relations  to  the  illustration, 
we  have  the  proportion:  — 

775  :  760  :  :  x :  10;    x  =  10.197  liters. 

The  mathematical  operation  of"  finding  the  volume  a  gas  would 
occupy  at  760  mm.  is  called  reducing  to  standard  pressure  or  cor- 
recting for  pressure. 

32.  Behavior    of    Gas    Volumes    under    Simultaneous 
Changes  in  Temperature  and  Pressure.  —  It  is  imma- 
terial whether  a  gas  is  subjected  to  changes  in  tempera- 
ture and  then  to  changes  in  pressure  or  to  both  at  once. 
Heat  and  pressure  act  independently.     Hence,  a  gas  vol- 
ume can  be  corrected  for  temperature  and  pressure  by  a 
single  calculation.     Thus,  if  the  observed  volume  is  10 
liters,  the  temperature  15°  C.,  and  the  pressure  775  mm., 
the  condensed  formula  is :  — 

X  =  10  X  2737(273  +  15)  X  775/76°;  x  =  9-666  liters 

Reduction  of  gas  volumes  to  standard  conditions  may  be  accom- 
plished by  substituting  the  observed  values  in  the  following  equation 
and  solving  for  V. 

V'P' 

~  760  (i  +  (.00366  X  t)  ) 

In  this  equation,  V  =  corrected  volume,  V  =  observed  volume, 
P'  =  observed  pressure,  t  =  observed  temperature.  (The  method 
of  deriving  this  equation  is  given  in  the  author's  Experimental 
Chemistry,  pp.  361-363.) 

33.  Weight  of  a  Liter  of  Oxygen  Gas. —  As  already 
stated  (12),  the  weight  of  a  liter  of  oxygen  gas  is  1.429 
gm.  at  o°   C.  and  760  mm.     This  value  (1.429)  is  found 
by  an  experiment  involving  several  steps,     (a)  Oxygen  is 


34  CHEMISTRY 

generated  from  a  mixture  of  potassium  chlorate  and  man- 
ganese dioxide,  and  the  weight  liberated  is  found  by  sub- 
tracting the  weight  of  the  oxygen  generator  after  the 
experiment  from  its  original  weight;  suppose  the  weight 
of  liberated  oxygen  is  2.322  gm.  (b)  The  oxygen  is  col- 
lected, its  volume  noted,  and  the  temperature  and  pres- 
sure also  read;  suppose  the  volume  of  oxygen  is  1.75  1., 
the  temperature  19°  C.,  and  the  pressure  755  mm.  (c) 
The  observed  volume  is  reduced  to  the  volume  at  standard 
conditions,  thus:  —  1.75  X  755/760  X  273/293  =  1.625. 
(d)  The  weight  of  one  liter  at  o°  C.  and  760  mm.  is  then 
found  to  be  1.429  gm.  by  dividing  2.315  (the  weight  of 
the  oxygen)  by  1.625  (the  corrected  volume  of  the 
oxygen),  thus:  —  2.322  -r-  1.625  =  1429.  (See  Part  II, 
Exp.  13.) 

EXERCISES 

1.  State  Boyle's  law.     Illustrate  it. 

2.  State  Charles's  law.     Illustrate  it. 

3.  Give  examples  of  (a)  expansion  and  of  (6)  contraction  of  gases  caused 
by  change  of  temperature. 

4.  Apply  Exercise  3  to  change  of  pressure. 

6.  Describe  (a)  a  centigrade  thermometer  and  (6)  a  barometer.  (For 
(a)  see  App.,  §  2.) 

6.  Explain  the  expression  "  a  gas  is  under  standard  conditions." 

7.  When  a  given  volume  of  gas  is  reduced  to  standard  conditions,  is  the 
weight  of  the  gas  changed?     Is  the  gas  itself  reduced  to  o°  C.  and  760  mm.? 

8.  From  the  data  given  in  33,  calculate  the  weight  of  a  liter  of  oxygen 
without  making  the  correction  for  temperature  and  pressure.     Compare  the 
result  with  the  correct  weight. 

9.  Write  a  brief  essay  on  Boyle's  contributions  to  science. 

PROBLEMS 

1.  Reduce  the  following  to  the  volume  occupied  at  760  mm. :  (a)  20  cc. 
at  745  mm.;  (&)  45  cc.  at  765  mm.;  (c)  450  cc.  at  755  mm.;  (d)  1.5  1.  at 
763  mm.;  (e)  2.5  1.  at  745  mm.;  (/)  500  cc.  at  75  cm.;  (g)  76  cc.  at  76  cm.; 
(ti)  900 1.  at  749  mm. 


SOME  PROPERTIES  OF  GASES  35 

2.  Reduce  the  following  to  the  volume  occupied  at  o°  C.:  (a)  170  cc. 
at  80°  C.;    (6)  450  cc.  at  15°  C.;    (c)  70.6  cc.  at  17°  C.;    (d)  49  cc.  at  19° 
C.;  («)  356  cc.  at  34°  C.;    (/)  48  cc.  at  27°  C. 

3.  Reduce  the  following  to  the  volume  at  standard  conditions:    (a) 
250  cc.  at  780  mm.  and  20°  C.;   (6)  140  cc.  at  745  mm.  and  21°  C. 

4.  Reduce  the  following  to  the  volume  at  standard  conditions:    (a) 
247  cc.  at  720  mm.  and  14°  C.;    (b}  1000  cc.  at  750  mm.  and  18°  C.;    (c) 
1480  cc.  at  765  mm.  and  81°  C. 

6.  A  volume  of  oxygen  measured  375  cc.  when  the  barometer  stood  at 
740  mm.,  and  its  temperature  was  27°  C.  What  would  be  the  volume  at 
the  standard  pressure  and  temperature? 

6.  Find  the  weight  of  29  cc.  of  oxygen  at  23°  C.  and  776  mm. 

7.  Calculate  the  weight  of  hydrogen  in  a  vessel  of  10  liters  capacity, 
filled  when  the  barometer  reads  756  mm.  and  the  thermometer  18°  C. 

8.  What  will  be  the  volume  of  100  cc.  of  hydrogen,  measured  at  14°  C. 
and  755  mm.,  when  it  is  reduced  to  o°  C.  and  760  mm.? 

9.  A  gas  measures  637  cc.  at  755  mm.  and  17°  C.      Find  its  volume 
at  760  mm.  and  o°  C. 

10.  A  gas  measures  100  liters  at  14°  C.  and  750  mm.     Find  its  volume 
at  standard  conditions. 

11.  Reduce  250  cc.  of  oxygen  gas  at  18°  C.  and  745  mm.  to  the  volume 
occupied  at  o°  C.  and  760  mm. 

12.  Reduce  189  cc.  of  hydrogen  at  15°  C.  and  750  mm.  to  the  volume  at 
o°  C.  and  760  mm. 


CHAPTER  V 

PROPERTIES    OF    WATER 

34.  Occurrence  in  Nature.  —  Water  is  always  present 
in  the  atmosphere  as  vapor.     In  the  liquid  state  water 
occurs  in  vast  quatities.     The  soil  and  plants  contain  con- 
siderable.    Many  common  foods  consist  largely  of  water 
(see  tables,  Chapter  XVII).     The  human  body  is  nearly 
70  per  cent  water.     (See  Part  II,  Exps.  14,  25.) 

35.  Natural  Waters.  —  Water  is  never  found  pure  in 
nature.     Even  rain  water,  which  is  usually    regarded  as 
the  purest  natural  water,  contains  material  washed  from 
the  atmosphere.     Surface  and  underground  water  dissolves 
substances  from  the  rocks  and  soil.     Water   containing 
certain  calcium  and  magnesium  compounds  is  hard,  but 
in  soft  water,  such  as  rain  water,  these  compounds  are 
absent  (249,  423). 

36.  River  water  contains  the  impurities  brought  by 
underground  and  surface  water;  it  is  also  often  polluted 
by  decaying  animal  and  vegetable  matter  or  by  refuse 
from  manufactories.     Ocean  water  contains  a  large  pro- 
portion of  common  salt.     Other  substances  are  present, 
especially  compounds  of  magnesium  and  calcium.     The 
peculiar  taste  of  ocean  water  is  due  to  the  presence  of 
these  substances. 

37.  Drinking    Water.  —  The  water  of  many  cities  is 
purified  by  filtering  it  on  a  large  scale  through  layers  of 


PROPERTIES  OF  WATER 


37 


sand  and  gravel  (Fig.  9) .  Such  a  filter  removes  bacteria 
almost  completely,  though  it  must  be  frequently  cleaned. 
Sometimes  the  water  is  stored  in  a  large  settling  basin 
or  tank  and  purified  before  filtration  by  adding  alum  or  a 
similar  substance,  which  causes 
the  suspended  matter  to  settle. 
Ozone  (20)  and  bleaching  powder 
(78)  are  used  as  purifiers  in  some 
localities.  (See  Part  II,  Exp.  207.) 


Fig.  9.  —  Section  of  a  Sand 
Filter  showing  Impure  Water 
(top),  Sand,  Gravel,  and 
Filtered  Water. 


38.  The  purity  of  drinking  water 
is  found  by  a  chemical  and  micro- 
scopic examination  of  a  sample,  sup- 
plemented by  a  rigid  sanitary  inspection  of  the  source  of  supply. 

Water  can  be  purified  by  distillation.  This  operation  is  often 
performed  in  the  laboratory  in  a  condenser,  which  is  shown  in  Fig. 
10  arranged  for  use.  The  condenser  consists  of  an  outer  tube 
A  A',  provided  with  an  inlet  and  an  outlet  for  a  current  of  cold  water, 


Fig.    10.  —  Condenser. 

which  surrounds  the  inner  tube  B  B'.  The  vapor  from  the  water 
boiling  in  the  flask  C  condenses  in  the  inner  tube,  owing  to  the 
decrease  in  temperature,  and  drops  off  the  lower  end  of  this  tube, 
as  the  distillate,  into  the  receiver  D,  while  the  impurities  remain 
behind  in  the  flask.  (See  Part  II,  Exp.  26.)  Other  forms  of 


38  CHEMISTRY 

condensers  are  used  (Fig.  n).  Cold  water  enters  the  condenser 
at  the  lower  inlet  and  is  kept  level  in  the 
chamber  by  the  upper  outlet.  The  chamber 
is  heated,  and  the  steam  in  passing  down 
through  the  condenser  drops  off  the  lower 
end  as  distilled  water.  Distilled  water  is 
prepared  on  a  large  scale  by  boiling  the 
water  in  a  metal  vessel  and  condensing  the 
vapor  in  a  block  tin  pipe  coiled  around  the 

inside    of   a   vessel  

through  which  a  cur- 
rent of  cold  water  is 
flowing  (Fig.  12).  Dis- 
tilled water  is  used  in 
the  chemical  labora- 
tory to  prepare  many 
solutions. 


Fig.   ii. —  Apparatus 
for  Distilling  Water. 


Fig.  12.  —  Coiled  Pipe 
Condenser. 


39.  Physical  Properties  of  Water. - 
At  ordinary  temperatures  pure  water 
is  a  tasteless  and  odorless  liquid.  It  is  usually  colorless, 
but  thick  layers  are  bluish.  Water  is  a  poor  conductor  of 
heat.  (See  Part  II,  Exp.  15.) 

Water  solidifies  or  freezes  at  o°  C.  (or  32°  Fahrenheit). 
When  water  freezes,  it  expands  about  one  tenth  of  its 
volume.  Hence  ice  floats.  The  specific  gravity  of  ice  is 
about  0.92. 

The  pressure  exerted  by  water  when  it  freezes  is  powerful. 
Vessels  or  pipes  completely  filled  with  water  often  burst  when  the 
water  freezes.  It  is  an  erroneous  but  popular  idea  that  "thawing 
out"  a  pipe  bursts  it.  As  a  matter  of  fact,  ice  contracts  when  it 
melts.  Pipes  crack  as  soon  as  the  water  freezes,  and  when  the  ice 
melts,  a  channel  is  left  for  the  water  to  flow  out  of  the  pipe. 


Ice  melts  at  o°  C.  (32°  F.),  which  is  also  the  freezing 
point  of  water.    Ice  often  crystallizes  in  forming,  but  in- 


PROPERTIES  OF  WATER 


39 


dividual  crystals  are  seldom  visible  except  during  the  first 
stages  of  the  process.  Snow  crystals  are  common  (Fig. 
13).  They  are  al- 
ways six-sided  or  six- 
pointed,  and  are 
formed  in  the  at- 
mosphere by  the 
freezing  of  water 
vapor. 

40.  Vapor  Pres- 
sure. —  Water  evap- 
orates at  all  temperatures.  If  water  is  heated  in  an  open 
vessel,  the  vapor  escapes  rapidly  until  the  thermometer 
reaches  100°  C.  (or  212°  F.).  At  this  point  water  boils,  i.e. 
it  changes  rapidly  into  vapor  without  rise  of  temperature. 
This  vapor,  if  allowed  to  escape  into  the  atmosphere,  cools 
and  condenses  quickly  into  a  cloud  of  minute  drops  of 
water.  This  cloud  is  popularly  called  steam. 


Fig.  13.  —  Snow  Crystals.    (From  a   photo- 
graph by  Wilson  A.  Bentley.) 


Fig.  14.  —  Experiment  to  Illustrate  Vapor  Pressure. 

The  water  vapor  that  escapes  from  the  surface  of  water 
produces  pressure,  which  is  called  vapor  pressure.  This  fact 
may  be  illustrated  by  the  apparatus  shown  in  Fig.  14.  If  a 


CHEMISTRY 


little  water  is  introduced  into  the  dry  bottle  (left)  by  pushing 
down  the  tube  that  contains  water  in  the  bulb,  the  colored 
liquid  in  the  U-shaped  tube  will  indicate  a 
pressure  inside  the  bottle  (right).  The  pres- 
sure of  water  vapor  depends  solely  on  the  tem- 
perature. This  is  readily  seen  by  comparing 
the  heights  of  the  mercury  in  the  barometer 
tubes  shown  in  Fig.  15.  In  the  tube  A  there 
is  no  water  vapor  in  the  space  above  the 
mercury,  and  the  height  of  the  mercury  is 
760  mm.  In  B  the  space  above  the  mercury 
is  filled  with  water  vapor  at  20°  C. ;  the  vapor 
exerts  a  pressure  and  forces  the  mercury  down 
to  nearly  742  mm.  That  is,  water  vapor  at 
20°  C.  exerts  a  pressure  equal  to  about  18 
mm.  of  mercury.  Similarly,  in  C  the  space 

Fig  I5 Experi-    *s  filled  with  water  vapor  at  50°  C.  and  the 

ment  to  show  mercury  is  forced  down  to  678  mm.,  the  water 
the  Relation  vapor  exerting  a  pressure  of  about  82  mm.  If 
between  Vapor  ^ne  vapor  Were  at  100°  C.  the  vapor  pressure 

would  be  760  mm.      The  latter  value  is   in- 
Temperature.  ' 

structive,  for  it  means  that  at  the  boiling 
point  of  water  (100°  C.)  the  vapor  pressure  just  balances  the 
normal  and  opposing  atmospheric  pressure.  The  pressure  ex- 
erted by  water  vapor  depends  solely,  as  shown  above,  on  the 
temperature  of  the  evaporating  water,  and  has  a  maximum 
value  for  each  temperature.  These  values  have  been  carefully 
determined  by  experiment,  and  can  be  found  in  the  Table  of 
Vapor  Pressure  given  in  the  Appendix,  §  4. 

A  practical  application  of  vapor  pressure  is  made  in  finding  accu- 
rately the  weight  of  a  liter  of  oxygen  and  in  similar  experiments 
where  gases  are  measured  over  water.  The  oxygen  gas  is  collected 
in  a  bottle  or  graduated  tube  inverted  in  a  vessel  of  water.  If  the 
gas  is  allowed  to  stand  confined  over  the  water  long  enough,  it 
becomes  saturated  with  water  vapor;  i.e.  the  tube  finally  contains 
a  mixture  of  oxygen  and  the  maximum  amount  of  water  vapor  at 


PROPERTIES  OF  WATER  41 

the  given  temperature.  In  such  a  mixture,  each  gas  shares  the 
total  atmospheric  pressure.  Hence  the  actual  pressure  exerted 
by  the  oxygen  is  found  by  subtracting  the  pressure  of  the  water 
vapor  from  the  total  pressure  (indicated  by  the  barometer).  The 
latter  method  is  used,  because  the  pressure  of  water  vapor  at  any 
temperature  is  known  and  can  be  taken  directly  from  the  table. 
Incorporating  this  fact  into  the  formula  given  in  Chapter  IV  for 
reducing  the  volume  of  a  gas  to  its  volume  at  o°  C.  and  760  mm. 
the  formula  becomes  — 

v  V  (P'  -  a) 


760(1+ (.00366X1)) 

In  this  formula  V  means  the  volume  of  dry  oxygen  at  o°  C.  and 
760  mm.,  and  a  means  the  vapor  pressure  (found  in  the  table  in  the 
Appendix,  §  4). 

41.  Chemical  Properties  of  Water.  —  Water  at  ordi- 
nary temperatures  interacts  with  certain  metals,  espe- 
cially calcium,  sodium,  and  potassium  (21).  Magnesium 
and  zinc  interact  with  boiling  water,  and  iron  with  steam 
(Fig.  6).  Water  is  decomposed  to  some  extent  into  its 
component  elements  (oxygen  and  hydrogen)  by  intense 
heat;  at  about  2000°  C.  the  decomposition  is  less  than 
2  per  cent.  As  the  temperature  falls,  the  elements  recom- 
bine  to  form  water.  Water  combines  directly  with  many 
oxides.  Thus,  lime,  which  is  calcium  oxide,  combines  di- 
rectly with  water  and  forms  a  compound  called  calcium 
hydroxide;  this  chemical  change  is  attended  by  consider- 
able heat.  Similarly,  sulphur  dioxide  forms  sulphurous 
acid.  These  chemical  changes  may  be  represented  thus :  — 

Calcium  Oxide  +  Water  =  Calcium  Hydroxide 

(Calcium-Oxygen)  (Calcium-Hydrogen-Oxygen) 

Sulphur  Dioxide  +  Water  =  Sulphurous  Acid 

(S  ulphur-Oxygen)  (Sulphur-Hydrogen-Oxygen) 


42  CHEMISTRY 

Such  oxides  are  sometimes  called  anhydrides.  Water  com- 
bines with  certain  solids  when  they  separate  from  a  solu- 
tion by  crystallization.  Thus,  from  a  solution  of  copper 
sulphate  blue  crystals  are  obtained,  which  when  heated 
give  off  water  and  crumble  to  a  gray  white  powder  (see 
Water  of  Crystallization,  49). 

42.  Solvent  Power  of  Water.  --  This  is  one  of  the  most 
conspicuous  properties  of  water.    Only  the  simpler  aspects 
are  treated  in  this  chapter.     More  extended  discussion 
may  be  found  in  Chapter  XIV. 

The  clear,  transparent  liquid  containing  a  dissolved  sub- 
stance is  called  a  solution.  The  liquid  in  which  the  sub- 
stance dissolves  is  called  the  solvent,  and  the  dissolved 
substance  is  called  the  solute.  Substances  differ  widely 
in  their  solubility.  A  solution  which  contains  a  small  pro- 
portion of  solute  is  called  a  dilute  solution;  one  contain- 
ing a  large  proportion  of  solute  is  called  a  concentrated 
solution.  Sometimes  the  terms  weak  and  strong  are  used 
instead  of  dilute  and  concentrated.  Other  solubility 
terms  are  defined  below  (45). 

43.  Solutions  of  Gases.  —  Water  dissolves  many  gases. 
The  solubility  varies  widely.     Some,  like  ammonia,  are 
very  soluble,  while  others,  such  'as  oxygen  and  hydrogen, 
are  only  slightly  soluble.    As  a  rule,  the  solubility  of  a  gas 
decreases  with  rise  of  temperature.     Pressure   influences 
the  solubility  of  gases.    Thus,  large  quantities  of  carbon 
dioxide  gas  are  forced  into  cylinders  full  of  water  in  pre- 
paring soda  water.     When  the  pressure  is  decreased  by 
opening  the  valve,  the  gas  escapes  rapidly  and  causes  the 
soda  water  to  froth  or  foam;   bubbles  caused  by  escaping 
carbon  dioxide  may  also  be  seen  when  the  stopper  is  re- 
moved from  a  bottle  containing  a  charged  beverage.    This 
rapid  escape  of  gas  is  called  effervescence.    Underground 


PROPERTIES  OF  WATER  43 

waters  often  contain  large  amounts  of  gases,  especially 
carbon  dioxide,  owing  to  the  great  pressure  to  which 
subterranean  gases  are  subjected.  Hence,  natural  mineral 
waters  often  effervesce  when  they  come  to  the  surface. 
(See  Part  II,  Exp.  17.) 

44.  Solutions    of    Liquids.  —  Some    liquids,    such    as 
alcohol  and  glycerin,  dissolve  in  water  in  all  proportions; 
others,  e.g.  kerosene  and  carbon  disulphide,  are  practically 
insoluble,  as  is  shown  by  the  fact  that  after  agitation  with 
water  they  separate  almost  entirely  as  distinct  layers. 
The  formation  of  separate  layers  must  not  be  accepted 
as  final  evidence  of  the  insolubility  of  the  liquid.     Ether 
and  water  form  two  layers,  but  each  dissolves  appreciably 
in  the  other.     The  upper  layer  consists  of  ether  and  a 
little  water;   the  lower  layer  is  the  opposite.    Alcohol  and 
water  form  no  such  layers,  not  simply  because  each  is 
soluble  in  the  other,  but  because  each  is  soluble  without 
limit  in  the  other;    i.e.  it  is  a  case  of  perfect  mutual 
solubility.     (See  Part  II,  Exp.  18.) 

45.  Solutions  of  Solids.  —  Water  dissolves  many  solids, 
and  such  solutions  are  very  useful.     The  solubility  of 
solids  in  water  depends  on  the  substance  itself  and  the 
temperature  of  the  water.     In  most  cases  solubility  in- 
creases with  a  rise  of  temperature;    hence  the  common 
practice  of  heating  to  hasten  solution.    A  few  solids  (e.g. 
calcium  hydroxide)  are  less  soluble  in  hot  water  than  in 
cold,  and  a  few  others  (e.g.  sodium  chloride)  dissolve  to 
about  the  same  degree  in  hot  and  cold  water.     A  given 
weight  of  water  at  a  fixed  temperature  will  dissolve  only  a 
definite  weight  of  solid;  and  this  is  the  case,  even  though 
more  undissolved  solid  is  available  for  solution.     A  solu- 
tion conforming  to  the  conditions  just  stated  is  said  to 
be  saturated.        For  general  purposes,  solubility  may  be 


44  CHEMISTRY 

TABLE  OF  THE  SOLUBILITY  OF  SOLIDS  IN  WATER 


Number  of  Grams  in  Solution 

Solids 

in  100  Grams  of  Water 

20°  C. 

100°  C. 

Calcium  Chloride 

74-5 

159.0 

Calcium  Hydroxide 

•  165 

.077 

Magnesium  Sulphate 

36-2 

73-8 

Potassium  Bichromate 

13.0 

IO2.O 

Potassium  Nitrate 

31-6 

246.0 

Sodium  Chloride 

36.0 

39-8 

expressed  by  such  terms  as  insoluble,  slightly  soluble,  or 
very  soluble.  It  is  more  accurate  to  represent  the  amount 
of  solvent  by  100  gm.;  on  this  basis  the  solubility  of  a 
solid  becomes  the  number  of  grams  of  solid  dissolved  by 
100  gm.  of  water.  (See  Part  II,  Exps.  19,  27.) 

46.  Solubility  Curves.  —  The  table  of  solubilities  just  given  is 
limited  to  two  temperatures.  A  more  complete  way  of  represent- 
ing the  solubility  of  a  substance  is  by  a  solubility  curve.  The 
curves  of  several  substances  are  shown  in  Fig.  16.  The  tempera- 
ture is  read  from  the  vertical  lines  and  the  number  of  grams  of 
solute  in  100  gm.  of  water  from  the  horizontal  lines.  For  example, 
if  we  wish  to  know  the  temperature  at  which  40  gm.  of  potassium 
chlorate  are  held  in  solution  by  100  gm.  of  water,  it  is  only  neces- 
sary to  find  where  the  horizontal  line  numbered  40  cuts  the  potas- 
sium chlorate  curve,  and  then  follow  the  vertical  line  down  to  the 
temperature  number,  where  80  C.  is  found. 

47.  Solution  and  Crystallization.  —  If  hot  solutions 
are  cooled  or  concentrated  solutions  are  evaporated,  the 
solute  separates  from  the  solvent  in  crystals;  the  process 
of  obtaining  them  is  called  crystallization.  The  shape 
and  color  of  the  crystals  are  characteristic  of  the  particu- 
lar substance  and  serve  to  identify  it.  .Thus,  common 
salt  crystallizes  in  white  cubes.  (See  Part  II,  Exp.  20.) 


PROPERTIES  OF  WATER 


45 


48.  Supersaturated  Solutions.  —  Crystals  are  not  al- 
ways deposited  as  just  stated.  Thus,  a  hot,  very  con- 
centrated solution  of  some  solids,  such  as  sodium  sulphate 

150 


0°    0°   20°   30'   40°   50°   60°   70°   80°   90°   100° 

Temperature 
Fig.  1 6. — Solubility  Curves. 

and  sodium  thiosulphate,  deposits  no  crystals  when  the 
clear  solution  cools.  Solutions  which  contain  more  solute 
than  is  needed  for  normal  saturation  are  called  super- 
saturated. Supersaturation  can  occur  only  when  the 
undissolved  solid  is  absent.  If  a  fragment  of  the  solid 
is  dropped  into  the  supersaturated  solution,  crystals  very 


46  CHEMISTRY 

soon   begin   to  form  upon  the  fragment.     (See  Part  II, 
Exp.  28.) 

49.  Solution  and  Water  of  Crystallization.  —  Crystals 
of  some  substances  deposited  from  solutions  contain  water 
which  is  an  essential  part  of  the  compound.    The  combined 
water  must  not  be  confused  with  water  which  adheres  to 
a  crystal  or  is  inclosed  in  it.     Crystals  containing  com- 
bined water  are  dry  even  after  the  crystals  are  powdered. 
The  combined  water  can  be  removed  by  heat  or  some- 
times merely  by  exposure  to  air.     Loss  of  water  is  usually 
attended  by  loss  of  color  and  always  by  loss  of  crystalline 
appearance.     Thus,  blue  crystallized  copper  sulphate  loses 
color  slowly  at  ordinary  temperatures  and  very  rapidly 
when  heated,  finally  becoming  a  gray  powder.    The  pro- 
portion of  combined  water  in  crystals  is  constant  in  the 
same  compound,  but  in  different  substances  the  proportion 
varies  between  wide  limits.     Water  chemically  combined 
in  a  crystal  and  readily  removed  in  a  definite  proportion 
by  heating  is  called  water  of  crystallization.     Compounds 
containing  water  of  crystallization  are   sometimes  called 
hydrates  or  hydrated  compounds.    Conversely,  compounds 
which  have  been  deprived  of  water  of  crystallization  are 
said  to  be  anhydrous  or  dehydrated.     For  example,  blue 
crystallized  copper  sulphate  is  a  hydrate  of  the  compound 
copper  sulphate,  but  after  the  blue  compound  has  been 
heated,  it  becomes  anhydrous  or  dehydrated  copper  sul- 
phate,  which  is  a  gray  powder.    Anhydrous  compounds 
often  readily  become  hydrated  again.     Thus,  when   the 
gray  anhydrous  copper  sulphate  is    added  to   water,    a 
blue  solution  is  produced  from  which  blue  crystals  of  hy- 
drated copper  sulphate  are  readily  obtained.     (See  Part 
II,  Exps.  21,  22.) 

50.  Efflorescence.  —  Some  substances  lose  their  water 


PROPERTIES  OF  WATER 


47 


of  crystallization  wholly  or  in  part  by  exposure  to  air. 
This  property  is  called  efflorescence,  and  such  substances 
are  said  to  be  efflorescent  or  to  effloresce.  Crystals  of 
washing  soda,  alum,  and  borax  effloresce  readily.  (See 
Part  II,  Exp.  23.) 

An  explanation  of  efflorescence  is  found  in  the  principle  of  vapor 
pressure.  Substances  containing  water  of  crystallization  exert  a 
vapor  pressure.  If  this  vapor  pressure  is  greater  than  the  pressure 
of  the  water  vapor  in  the  atmosphere,  the  substance  loses  water 
until  the  vapor  pressures  are  equal  or  until  all  the  water  has  escaped 
from  the  substance. 

51.  Deliquescence.  —  Many  substances  absorb  water 
when  exposed  to  air,  become  moist,  and  sometimes  even 
dissolve  in  the  absorbed  water.  Cal- 
cium chloride,  potassium  carbonate, 
zinc  chloride,  sodium  hydroxide,  mag- 
nesium chloride,  and  potassium  hy- 
droxide belong  to  this  class.  This 
property  is  called  deliquescence  and 
the  substances  are  said  to  deliquesce, 
or  to  be  deliquescent.  Deliquescence 
is  a  property  of  very  soluble  sub- 
stances. (See  Part  II,  Exp.  24.)  Fig.  17.- A  Desiccator. 

Deliquescence  can  be  explained  thus:  Water  vapor  from  the  air 
condenses  on  the  surface  of  the  solid  and  produces  a  very  concen- 
trated solution,  which  has  a  vapor  pressure  much  lower  than  the 
average  pressure  of  the  water  vapor  in  the  air;  the  solution,  there- 
fore, continues  to  take  up  water  until  its  vapor  pressure  equals  the 
pressure  of  the  water  vapor  in  the  air.  Common  salt  often  de- 
liquesces, especially  in  damp  weather,  owing  to  small  quantities  of 
magnesium  and  calcium  chlorides  which  are  present  as  impurities. 
The  property  of  deliquescence  is  utilized  in  the  laboratory  to  dry 
substances,  calcium  chloride  being  often  employed  for  this  purpose. 
One  form  of  apparatus  used  is  called  a  desiccator  (Fig.  17). 


48  CHEMISTRY 


EXERCISES 

1.  Why  is  sea  water  salt? 

2.  Essay  topics:    (a)  Purification  of  drinking  water.     (b)  Distillation. 
(c)   Water  as  an  erosive  agent,    (d)  Crystals,     (e)   Vapor  pressure.     (/) 
Chemical  properties  of  water. 

3.  Define  and  illustrate  (a)  water  of  crystallization,  (b)  efflorescence, 
(c)  deliquescence,  (d)  anhydrous,  (e)  dehydrated. 

4.  Define  and  illustrate  (a)  solution,  (b)  solvent,  (c)  solute,  (d)  dilute, 
(e}    concentrated,   (/)   unsaturated  solution,    (g)   saturated   solution,   (h) 
supersaturated  solution,  (i)  solubility,  (j)  solubility  curve. 

5.  How  might  sea  water  be  rendered  suitable  for  drinking? 

6.  Practical  topics:    (i)  How  would  you  prove  a  liquid  is  pure  water? 
(2)  What  conditions  are  favorable  for  (a)  evaporation,  (b)  efflorescence, 
(c)  deliquescence?     (3)  Suggest  experiments  (a)  to  find  the  solubility  of 
a  solid  in  water  at  40°  C.,  (b)  to  show  that  water  from  a  crystal  is  not  water 
of  crystallization,  (c)  to  find  the  per  cent  of  water  in  a  potato. 

PROBLEMS 

1.  Find  the  volume  of  the  dry  gas  at  o°  C.  and  760  mm.  in:    (a)  80  cc. 
at  750  mm.  and  17°  C.;    (b)  80  cc.  at  745  mm.  and  19°  C.;    (c)    100  cc.  at 
765  mm.  and  17.5°  C.;   (d)  97  cc.  at  757  mm.  and  20.5°  C. 

2.  Plot  the  following  data  on  cross  section  paper  and  draw  the  solu- 
bility curve  of  the  substance:  Temperature  —  o,  10,  20,  30,  40,  50,  55;  cor- 
responding solubility  (i.e.  number  of  gm.  soluble  in  100  gm.  of  water)  — 
13,  21,  31,  45,  64,  86,  100. 

3.  If  the  density  of  ice  is  0.92,  what  volume  will  a  liter  of  water  at 
4°  C.  occupy  when  frozen? 

4.  By  use  of  the  solubility  curves  in  Fig.    16   answer   the   following: 
(a)  How  many  gm.  of  sodium  chloride  are  in  solution  at  20°,  30°,  55°,  65°, 
o°,  i oo°?     (6)  At  what  temperatures  are  60  gm.  and  95  gm.  of  potassium 
bromide  in  solution?     (c)  Compare  the  solubility  of  sodium  nitrate  and 
sodium  chloride.    How  much  of  each  is  in  solution  at  20°,  25°,  30°? 

6.  Calculate  the  per  cent  of  water  of  crystallization  in  each  crystal- 
lized substance  from  the  following:  (a)  5  gm.  of  aluminium  sulphate 
lose  2.43  gm.  on  heating;  (b)  7  gm.  of  calcium  sulphate  lose  1.464  gm.; 
(c)  3  gm.  of  cadmium  nitrate  lose  .7  gm.;  (d)  3  gm.  of  cobalt  nitrate  lose 
i.ii3gm. 


CHAPTER   VI 

COMPOSITION  OF  WATER  —  HYDROGEN  DIOXIDE 

Water  is  a  compound  of  hydrogen  and  oxygen.  That 
is,  its  constituents  are  the  elements  hydrogen  and  oxygen, 
and  they  are  chemically  combined  in  a  fixed  ratio. 

52.  The  Composition  of  a  Compound  is  determined  either 
by  analysis  or  synthesis,  i.e.  by  taking  it  apart,  directly 
or  indirectly,  or  by  putting  its  parts  together.    Sometimes 
both  methods  are  used. 

Composition  may  be  studied  qualitatively  and  quantitatively. 
A  qualitative  experiment  aims  to  discover  what  elements  or  groups 
of  elements  constitute  a  compound.  A  -quantitative  experiment  is 
an  accurate  determination  of  the  proportion,  by  weight  or  volume, 
of  the  constituents  of  a  compound. 

53.  Qualitative    Composition    of    Water.  —  The    fact 
that  water  contains  the  elements  hydrogen  and  oxygen 
can  be  shown  in  several  ways,    (i)  Metals  such  as  calcium, 
sodium,  and  iron  liberate  hydrogen  from  water  and  form 
simultaneously  compounds  containing  oxygen  (21).     (2) 
When  an  electric  current  is  passed  through  an  acid  solu- 
tion of  water,  hydrogen  and  oxygen  are  liberated  (54  i). 
(3)  Water  is  produced  when  hydrogen  is  burned  in  air  or 
in  oxygen  (24). 

The  fact  that  sodium  forms  an  oxygen  compound,  viz. 
sodium  hydroxide,  by  interaction  with  water  can  be  readily 
shown  by  experiment.  If  red  litmus  paper  is  put  into 
the  water  from  which  the  sodium  has  liberated  hydrogen, 
the  litmus  paper  becomes  blue.  This  change  of  color 


CHEMISTRY 


from  red  to  blue  shows  that  a  base  is  in  the  water,  because 
bases  turn  red  litmus  paper  blue.  The  base  can  be  ob- 
tained as  a  white  solid  by  evaporating  the  water.  If  a 
little  of  the  solid  is  heated  in  the  flame,  the  persistent, 
intense  yellow  color  imparted  to  the  flame  proves  that 
the  white  solid  is  a  sodium  compound. 
Further  tests  could  be  applied  to  show 
that  it  belongs  to  the  class  of  compounds 
called  hydroxides.  Sodium  hydroxide  is 
a  compound  of  sodium,  hydrogen,  and 
oxygen,  and  is  formed  by  replacing  half 
of  the  hydrogen  of  water  by  sodium. 

A  simple  experiment  shows  directly 
that  oxygen  is  a  chemical  constituent  of 
water,  viz.  the  exposure  of  chlorine  water 
to  sunlight.  (Chlorine  water  is  prepared 
by  saturating  water  with  chlorine  gas  — 
an  element  to  be  studied  in  Chapter  IX.) 
A  tube  about  a  meter  long  and  closed  at 
one  end  is  completely  filled  with  chlorine 
water,  the  open  end  is  immersed  in  a 
vessel  containing  some  of  the  same  solu- 
tion, and  the  whole  apparatus  is  placed 
Fig.  1 8.— Preparing  in  the  direct  sunlight.  Bubbles  of  gas 
Oxygen  from  soon  appear  in  the  liquid,  and  after  a  few 

Chlorine  Water. 

hours  a  small  volume  of  gas  collects  at 
the  top  of  the  tube  (Fig.  18).  If  the  tube  is  closed  (by 
the  thumb  or  finger),  removed,  and  inverted,  the  gas  will 
rise  to  the  open  end,  where  it  can  be  shown  to  be  oxy- 
gen by  the  usual  test,  viz.  relighting  a  glowing  joss  stick 
or  splint  of  wood.  (See  Part  II,  Exp.  29.) 

54.   The    Quantitative   Composition   of   Water. --The 
quantitative  composition  of  water  is  obtained  by  a  de- 


COMPOSITION  OF  WATER 


51 


termination  of  its  volumetric  and  its  gravimetric  composi- 
tion, that  is,  the  proportion  in  which  hydrogen  and  oxy- 
gen are  united  in  water.  Volumetric  means  "  by  volume  " 
and  gravimetric  "by  weight." 

The  volumetric  composition  of  water  has 
been  determined  by  analysis  and  synthesis. 

i.  Analysis.  Reference  has  repeatedly 
been  made  to  the  fact  that  water  can  be 
decomposed  into  its  constituents  by  an 
electric  current.  The  decomposition  of 
water  by  electricity,  or,  as  it  is  called 
traditionally,  the  electrolysis  of  water,  is 
accomplished  in  a  special  form  of  apparatus 
(Fig.  19).  Pure  water  does  not  conduct 
electricity,  so  a  mixture  of  water  (10  vols.) 
and  concentrated  sulphuric  acid  (i  vol.)  is 
poured  into  the  apparatus  until  the  reservoir 
is  half  full  after  the  stopcocks  have  been 
closed.  As  soon  as  an  electric  battery  of 
three  or  more  cells  is  connected  with  the 

piece  of  platinum  near  the  bottom  of  each  Fig.  19. Appa- 

tube,  bubbles  of  gas  appear  on  the  platinum,  ratus  for  the 
rise,  collect  in  the  upper  part  of  the  tubes, 
and  slowly  force  the  liquid  from  each  tube 
into  the  reservoir.  The  volume  of  gas  is  greater  in  one 
tube.  Assuming  that  the  tubes  have  the  same  diameter, 
the  gas  volumes  are  in  the  same  ratio  as  their  heights, 
which  will  be  found  by  measurement  to  be  approximately 
two  to  one.  Tests  show  that  the  gas  having  the  larger 
volume  is  hydrogen  and  that  the  other  gas  is  oxygen. 
(See  Part  II,  Exp.  30.) 

Owing  to  slight  errors,  the  quantitative  result  obtained 
by  this  method  is  approximate. 


CHEMISTRY 


2.  Synthesis.  An  accurate  determination  can  be  made 
by  exploding  a  mixture  of  known  volumes  of  hydrogen 
and  oxygen  in  a  eudiometer. 

A  simple  sketch  of  a  convenient  form  of  apparatus  is 

shown  in  Fig.  20.  The 
essential  part  is  the 
eudiometer  F.  In  this 
glass  tube  the  gases  are 
accurately  measured 
and  exploded.  The  elec- 
tric spark  that  causes 
the  explosion  is  obtained 
from  an  induction  coil 
and  battery.  The  spark 
leaps  across  the  space 
between  the  platinum 
wires  at  the  top  of  the 
eudiometer,  and  the 
heat  produced  by  the 

Fig.  20.  -Apparatus  for  Determining  the    spark  causes  the  hydro. 
Volumetric  Composition  of  Water. 

gen  and  oxygen  to  com- 
bine and  form  water.  Omitting  details,  oxygen  and 
hydrogen  are  introduced  separately  into  the  eudiometer, 
measured,  and  exploded.  The  quantity  of  water  formed 
by  the  union  of  the  hydrogen  and  oxygen  is  too  minute 
to  measure.  The  most  accurate  experiments  give  the 
ratio  2.0027  ^0  i,  but  as  usually  stated  two  volumes  of 
hydrogen  combine  with  one  volume  of  oxygen  to  form 
water. 

The  gravimetric  composition  of  water  has  been  accu- 
rately determined  by  synthesis.  The  apparatus  is 
shown  in  Fig.  21.  It  was  first  weighed  vacuous  (i.e.  free 
from  air  or  other  gases) .  The  tubes  aa  were  then  connected 


COMPOSITION  OF  WATER 


S3 


with  the  weighed  reservoirs  of  oxygen  and  hydrogen,  and 
the  oxygen  was  introduced.  Sparks  were  next  passed 
between  the  platinum  wires  cc,  and  the  heat 
ignited  the  hydrogen,  which  was  slowly  ad- 
mitted, the  combination  of  the  gases  taking 
place  at  bb.  The  water  vapor  condensed  in 
the  tube  dd,  the  lower  portion  of  which  was 
immersed  in  cold  water.  The  combustion  of 
the  hydrogen  was  continued  until  a  suitable 
weight  of  water  was  formed.  The  water  and 
its  vapor  were  then  converted  into  ice  by  put- 
ting the.  apparatus  into  a  freezing  mixture; 
the  residual  mixture  of  gases  was  drawn  off 
and  analyzed,  passing  in  its  exit  through  tubes 
of  phosphorus  pentoxide  in  ee  which  retained 
all  traces  of  water.  The  whole  apparatus 
was  finally  weighed,  the  increase  being  the 
weight  of  the  water  formed  by  the  combina-  Morley'sAp- 
tion  of  known  weights- of  hydrogen  and  oxy-  .^e^^inin0r 
gen.  As  the  result  of  exceptionally  careful  the  Gravi- 
experiments  the  ratio  of  hydrogen  to  oxygen  by  metric  Corn- 
weight  was  found  to  be  i  to  7.9395.  The  com- 
plete  synthesis  was  made  in  1895  by  Morley. 

55.  Summary.  —  Experiment  shows  that  water  con- 
sists of  hydrogen  and  oxygen  combined  in  a  fixed  ratio  by 
weight,  viz.  i  to  7.9395;  they  are  also  combined  in  the 
ratio  of  2.0027  to  I  by  volume.  Usually  these  ratios  are 
stated  approximately  as  2  to  1 6  by  weight  and  2  to  i  by 
volume.  Sometimes  the  gravimetric  composition  of  water 
is  stated  in  per  cent,  the  values  being  11.18  per  cent 
hydrogen  and  88.82  per  cent  oxygen. 


Fig.  21. — 


of 


54  CHEMISTRY 

EXERCISES 

1.  Can  water  be  correctly  called  "hydrogen  oxide"  ?     Why? 

2.  What  does  Lavoisier's  experiment  (21)  show  about  the  composition 
of  water? 

3.  How  is  the  electrolysis  of  water  accomplished?    What  does  it  show 
about  the  composition  of  water? 

4.  Compare  the  volumetric  and  the  gravimetric  composition  of  water. 

5.  What  does  the  burning  of  hydrogen  prove  about  the  composition  of 
water? 

PROBLEMS 

1.  Suppose  15  gm.  of  water  are  decomposed.    What  weight  of  (a)  oxygen 
and  (b)  hydrogen  is  produced?    What  volume  (at  o°  C.  and  760  mm.)  of 
(c)  oxygen  and  (d)  hydrogen? 

2.  What  volume  of  oxygen  (at  o°  C.  and  760  mm.)  must  be  used  to  unite 
with  175  gm.  of  hydrogen  to  form  water? 

3.  What  volume  of  hydrogen  (at  o°  C.  and  760  mm.)  must  be  used  to 
convert  175  gm.  of  oxygen  into  water? 

4.  A  mixture  of  500  cc.  of  oxygen  and  1250  cc.  of  hydrogen  (both  at 
o°  C.  and  760  mm.)  is  exploded.    What  weight  of  water  is  formed? 

6.  50  cc.  of  oxygen  are  mixed  with  500  cc.  of  hydrogen,  both  measured 
at  the  normal  temperature  and  pressure.    An  electric  spark  is  passed  through 
the  mixture.    What  volume,  if  any,  of  gas  will  remain,  and  how  would  you 
ascertain  whether  it  is  hydrogen  or  oxygen? 

Hydrogen  Dioxide 

56.  Hydrogen  Dioxide  is  composed  of  hydrogen 
and  oxygen.  But  the  proportion  of  the  components 
is  not  the  same  as  in  water.  It  contains  approximately 
two  parts  of  hydrogen  and  thirty-two  of  oxygen  by  weight. 
It  is  often  called  hydrogen  peroxide.  The  ordinary  com- 
mercial solution  contains  about  three  per  cent  of  hydrogen 
dioxide.  It  has  a  sharp,  pungent  odor,  and  a  bitter,  me- 
tallic taste.  The  solution  is  somewhat  unstable,  and  de- 
composes slowly  into  water  and  oxygen. 

It  is  used  to  bleach  human  hair,  ostrich  feathers,  fur, 
silk,  wool,  cotton,  bone,  and  ivory,  and  as  an  antiseptic 
in  dentistry  and  surgery. 


CHAPTER  VII 

LAW  AND  THEORY  —  LAWS  OF  DEFINITE  AND  MULTIPLE 

PROPORTIONS  —  ATOMIC  THEORY  — ATOMS  AND 

MOLECULES  —  SYMBOLS  AND  FORMULAS 

57.  Law  and  Theory.  —  We  discover  facts  by  observa- 
tion and  experiment.    Facts  which  always  occur  under  the 
same  conditions  soon  become  well  established.     Related 
facts  are  often  summarized  in  a  brief  general  statement 
called  a  law.    The  explanation  we  give  of  facts,  especially 
groups  of  related  facts,  is  called  a  theory. 

Laws  and  theories  are  of  great  service  in  chemistry, 
since  they  help  us  gather  into  intelligible  statements  our 
knowledge  of  a  vast  number  of  related  facts.  They  also 
help  us  to  discover  new  facts  and  interpret  phenomena. 

58.  Law   of   the    Conservation   of   Matter.  —  Experi- 
ment shows  that  in  a  chemical  change  the  total  weight  of 
the  matter  involved  is  not  altered.    Substances  are  trans- 
formed, but  the  weight  of  the  substances  entering  into  a 
chemical   change   equals   the   weight   of    the   substances 
resulting   from   the   chemical   change.      This   feature  of 
chemical  change  may  be  summed  up  by  the  law  of  the 
conservation  of  matter,  which  is  preferably  stated  thus:  - 

No  weight  is  lost  or  gained  in  a  chemical  change. 

59.  Law  of  Constant  Composition.  —  It  has  been  shown 
by  experiment  that  water  contains  88.82  per  cent  oxygen 
and  11.18  per  cent  hydrogen  (55).     Similar  experiments 
show  that  in  all  chemical  compounds  the  different  con- 
stituents are  present  in  a  definite  and  constant  propor- 
tion by  weight.    This  general  fact  may  be  stated  in  the 


56  CHEMISTRY 

form  of  a  law,  called  the  Law  of  Constant  Composition 
or  the  Law  of  Definite  Proportions,  thus:  - 

A  given  chemical  compound  always  contains  the  same 
elements  in  the  same  proportions  by  weight. 

(See  Part  II,  Exp.  32.) 

This  law  is  one  of  the  fundamental  laws  of  chemistry.  It  was 
established  as  the  outcome  of  a  controversy  between  two  French 
chemists,  Proust  (1755-1826)  and  Berthollet  (1748-1822).  Subse- 
quent experiments  by  the  Belgian  chemist  Stas  (1813-1891)  and 
the  American  chemist  Richards  (1868-)  have  firmly  established  our 
belief  in  the  accuracy  of  this  law. 

60.  Law  of  Multiple  Proportions.  —  The  composition 
of  compounds  is  usually  expressed  in  per  cent;  but  in  the 
case  of  a  series  of  compounds  percentage  reveals  nothing 
about  multiple  relations.  If,  however,  a  fixed  weight  of 
one  constituent  is  adopted  as  a  basis  of  comparison,  and 
the  composition  of  the  series  of  compounds  is  expressed 
in  terms  of  this  weight,  then  the  simple  multiple  relation 
which  exists  between  the  weights  of  the  other  constituent 
(or  constituents)  may  be  clearly  seen.  Thus,  no  multiple 
relation  is  apparent  in  the  statement  that  the  two  com- 
pounds of  carbon  and  oxygen  contain  respectively  27.27 
and  42.85  per  cent  of  carbon  and  72.72  and  57.14  per  cent 
of  oxygen.  But  if  we  adopt  i  (or  any  other  number)  as 
the  weight  of  carbon  in  each  compound,  the  weights  of 
oxygen  will  be  in  the  simple  ratio  of  2  to  i ;  i.e.  the  weight 
of  oxygen  in  one  compound  is  a  simple  multiple  of  its 
weight  in  the  other.  Similar  results  may  be  worked  out 
with  other  series  of  compounds. 

The  general  fact  of  multiple  proportions  is  expressed  as 
the  Law  of  Multiple  Proportions,  thus:  - 

When  two  (or  more)  elements  unite  to  form  a  series  of 
compounds,  a  fixed  weight  of  one  element  always  combines 


LAW  AND  THEORY  57 

with  such  weights  of  the  other  element  (or  elements)  that  the 
ratio  between  these  different  weights  can  be  expressed  by  small 
whole  numbers. 

61.  The  Atomic  Theory.  —  The  theory  that  explains 
the  facts  summarized  in  the  laws  just  discussed  is  called 
the  atomic  theory.    It  was  proposed  by  Dalton,  an  Eng- 
lish chemist,  about  1805.     According  to  this  theory,  (i) 
an  element  is  made  up  of  a  vast  number  of  very  small 
particles  called  atoms;  (2)  atoms  of  the  same  element  have 
the  same  weight;    (3)  atoms  of  different  elements  differ 
from  each  other  in  weight;    (4)  chemical  change  is  the 
union,  separation,  or  exchange  of  undivided  atoms.    The 
atomic  theory  means  in  a  few  .words  that  matter  is  com- 
posed  of   atoms,   which   remain   undivided   in   chemical 
changes.     By  means  of  the  atomic  theory,  many  facts 
about  substances  and  chemical  change  can  be  made  very 
much  clearer. 

62.  Atoms  and  Molecules.  —  The  term  molecule  is  ap- 
plied to  particles  which  consist  of  two   or   more   atoms 
chemically  combined;    if  the  atoms  in  a  molecule  are 
alike,  the  molecule  is  a  molecule  of  an  element,  but  if  the 
atoms  are  different,  then  the  molecule  is  a  molecule  of  a 
compound.     The  term  atom  is  reserved  to  designate  the 
smallest  particle  of  an  element  that  participates  in  a 
chemical  change.     The  molecule  is  sometimes  spoken  of 
as  the  physical  unit,  because  in  most  physical  changes 
molecules  are  not  decomposed.     Whereas   the  atom  is 
called  the  chemical  unit,  because  it  is  the  part  of  a  mole- 
cule that,  as  a  rule,  is  transferred  unchanged  in  chemi- 
cal changes.     Molecules  will  be  discussed  again.     (See 
Chapter  XII). 

Although  the  atom  is  conceived  to  pass  as  a  whole  from  com- 
pound to  compound,  it  should  not  be  inferred  that  atoms  do  not 


58  CHEMISTRY 

decompose  under  any  conditions.  The  phenomena  exhibited  by 
compounds  of  radium  show  that  there  are  particles  smaller  than 
atoms.  (See  Radioactivity.)  These  very  small  particles  are  called 
corpuscles  or  electrons.  However,  the  atom  is  the  chemical  unit, 
and  whether  or  not  it  is  a  complex  group  of  smaller  individuals,  its 
weight  is  not  altered  in  chemical  changes. 

63.   Interpretation  of  Laws  by  the  Atomic  Theory.  - 

First,  let  us  picture  a  chemical  change  in  terms  of  the 
atomic  theory,  e.g.  the  combination  of  copper  and  oxygen. 
An  appreciable  mass  of  copper  consists  of  many  millions 
of  atoms  of  copper;  a  mass  of  oxygen  likewise  consists  of 
atoms  of  oxygen.  When  the  chemical  change  occurs 
between  copper  and  oxygen,  atoms  of  copper  combine 
with  atoms  of  oxygen  and  form  molecules  of  a  compound 
called  copper  oxide.  And  this  combining  of  atoms  into 
molecules  continues  until  the  atoms  of  copper  or  of  the 
oxygen  (or  under  certain  conditions  the  atoms  of  both 
substances)  have  been  used  up.  Furthermore,  this  chem- 
ical change  takes  place  not  only  between  vast  numbers  of 
atoms,  but  the  quantitative  aspects  of  this  multitude  of 
changes  conform  to  the  atomic  theory.  This  latter  point 
needs  explanation,  because  it  emphasizes  the  chief  feature 
of  the  atomic  theory,  viz.  agreement  with  certain  funda- 
mental laws  of  chemical  change.  These  laws,  we  have 
already  found,  are  the  law  of  the  conservation  of  matter, 
the  law  of  constant  composition,  and  the  law  of  multiple 
proportions,  (i)  According  to  the  atomic  theory  the 
weight  of  an  atom  is  never  changed.  In  the  case  of  cop- 
per and  oxygen  the  weight  of  the  copper  oxide  formed 
equals  the  sum  of  the  weights  of  the  copper  and  oxygen 
used  up.  Inasmuch  as  all  other  chemical  changes  have 
this  characteristic,  viz.  unvarying  total  weight,  it  is  obvi- 
ous that  the  atomic  theory,  which  assumes  unchanging 


LAW  AND  THEORY  59 

weights  of  atoms,  is  in  accord  with  the  law  of  the  con- 
servation of  matter.  (2)  Again,  according  to  the  atomic 
theory,  when  copper  combines  with  oxygen,  molecules  of 
copper  oxide  are  formed  by  the  union  of  some  whole  num- 
ber of  atoms  of  copper  with  some  whole  number  of  atoms 
of  oxygen.  Each  molecule  of  copper  oxide  would  there- 
fore consist  of  one  or  more  atoms  of  copper  united  with 
one  or  more  atoms  of  oxygen,  and  the  composition  of  each 
molecule  of  copper  oxide  would  be  definite;  i.e'.  each  mol- 
ecule would  contain  the  same  elements  united  in  a  con- 
stant ratio  by  weight.  In  other  words,  copper  oxide 
would  always  be  found  to  consist  of  a  certain  per  cent  of 
copper  and  a  certain  per  cent  of  oxygen.  Since  all  other 
chemical  compounds  have  been  found  to  have  a  constant 
composition,  the  atomic  theory  harmonizes  with  the  law 
of  constant  composition.  (3)  Finally,  according  to  the 
atomic  theory  atoms  are  transferred  as  wholes;  this  means 
that  in  chemical  changes  there  are  no  fractions  of  atoms. 
Experiment  shows  that  there  are  two  oxides  of  copper. 
Each  contains  copper  and  oxygen  in  a  definite  ratio,  but 
the  ratios  are  different.  In  one  the  ratio  of  oxygen  to 
copper  is  1:4,  and  in  the  other  i  :8.  That  is,  the  weights 
of  copper  combined  with  a  fixed  weight  of  oxygen  are  in 
the  ratio  of  i :  2 ;  in  other  words,  if  a  molecule  of  one  com- 
pound consisted  of  one  atom  each  of  copper  and  oxygen, 
a  molecule  of  the  other  would  contain  two  of  copper  and 
one  of  oxygen.  Since  other  series  of  compounds  exhibit 
this  simple  multiple  relation,  it  is  evident  that  the  atomic 
theory  agrees  with  the  law  of  multiple  proportions. 

64.  Atomic  Weights.  —  According  to  the  atomic  theory 
atoms  of  the  same  element  always  have  the  same  weight 
but  atoms  of  different  elements  have  different  weights. 
This  means  (i)  that  an  atom  of  oxygen,  for  example, 


60  CHEMISTRY 

throughout  all  its  varied  changes  retains  its  weight,  and 
(2)  that  this  weight  differs  from  the  weight  of  other  kinds 
of  atoms.  The  weights  of  different  kinds  of  atoms  are 
called  the  atomic  weights  of  the  elements  or  briefly  atomic 
weights.  These  weights  have  been  determined  by  experi- 
ment and  a  table  giving  the  exact  and  approximate  values 
can  be  found  on  the  inside  of  the  back  cover  of  this  book. 
The  atomic  weights  are  relative  weights.  That  is,  the 
atomic  weight  of  copper  is  63.57,  not  63.57  gm.  or  any 
other  actual  weight,  but  63.57  as  l°ng  as  1 6  is  accepted  as 
the  standard  atomic  weight  of  oxygen. 

The  exact  determination  of  atomic  weights  is  a  difficult  task. 
Several  principles  must  be  considered  in  making  the  final  selection. 
Until  this  subject  is  discussed  (see  Chapter  XII),  it  will  be  well 
enough  to  regard  atomic  weights  as  the  numerical  values  of  the 
elements  in  chemical  changes  and  to  select  the  approximate  weights 
from  the  table  as  needed. 

65.  Chemical  Symbols  represent  single  atoms  of  the 
elements.     Thus,  H  represents  one  atom  of  hydrogen. 
If  more  than  one  uncombined  atom  is  to  be  designated, 
the  proper  numeral  is  placed  before  the  symbol,  thus:  - 
2H  means  2  atoms  of  hydrogen.     But  if  we  wish  to  rep- 
resent the  atoms  as  in  chemical  combination,  either  with 
themselves  or  with  other  atoms,  then  a  subscript  is  used 
instead  of  a  coefficient,  thus :  —  H2  means  2  atoms  of  hy- 
drogen in  combination,  as  in  H2O. 

Symbols  not  only  represent  atoms,  but  they  also 
express  atomic  weights.  Thus,  O  represents  one  atom 
of  oxygen  and  stands  for  the  atomic  weight  16. 

66.  Chemical  Formulas.  —  A  formula  is  a  group  of 
symbols  which  expresses  the  composition  of  a  compound. 
In  writing  a  formula,  the  symbols  of  the  atoms  making 
up  a  molecule  of  the  compound  are  placed  side  by  side. 


LAW  AND  THEORY  61 

Thus,  H2O  is  the  formula  of  water,  because  one  molecule 
consists  of  2  atoms  of  hydrogen  and  i  atom  of  oxygen. 
The  symbols  making  up  a  formula  might  be  written  in 
different  orders,  but  usage  has  determined  the  order  in 
most  cases.  A  formula  represents  one  molecule.  If  we 
wish  to  designate  several  molecules,  the  proper  numeral  is 
placed  before  the  formula,  thus:  —  2KC1O3  means  2  mole- 
cules of  potassium  chlorate.  In  certain  compounds  some 
of  the  atoms  act  like  a  single  atom  in  chemical  changes. 
This  is  often  expressed  by  inclosing  the  group  in  a  paren- 
thesis, e.g.  Zn(OH)Cl.  Sometimes  the  parenthesis  is 
replaced  by  a  period,  e.g.  C2H5.OH  and  CuSO4.5  H2O. 
The  period  and  parenthesis  are  occasionally  omitted, 
especially  if  the  composition  of  the  compound  is  well 
understood,  e.g.  NH4OH.  If  a  group  of  atoms  is  to  be 
multiplied,  it  is  placed  within  a  parenthesis.  Thus,  the 
formula  of  lead  nitrate  is  Pb(NO3)2.  This  means  that 
the  group  NO3  is  to  be  multiplied  by  2.  The  expression 
2  Pb(NO3)2  means  that  the  whole  formula  should  be 
multiplied  by  2.  Formulas  are  discussed  again  in  Chap- 
ter XII. 

67.  Molecular  Weights.  —  A  formula  represents  a 
molecular  weight  which  is  the  sum  of  the  atomic  weights 
represented  by  the  symbols  in  the  formula.  Thus,  the 
symbols  H  and  Cl  represent  1.008  and  35.46  respectively, 
and  the  formula  HC1  represents  1.008  +  35.46  or  36.468. 
If  we  know  the  formula  of  a  compound,  the  molecular 
weight  may  be  found  by  adding  the  atomic  weights  cor- 
responding to  the  atoms  in  the  formula.  Using  approxi- 
mate values,  the  molecular  weight  of  water  (H20)  is 
2  +  16  =  18;  the  weight  of  two  molecules  of  water  (2H2O) 
is  2(2  -f-  16)  =  36.  Similarly,  the  molecular  weight  of 
lead  nitrate  (Pb(NO3)2)  is  207  +  2(14  +  48)  =  331;  the 


62  CHEMISTRY 

weight  of  two  molecules  of  lead  nitrate  (2Pb(NO3)2)  is 
2  X  331  =  662. 

68.  Use  of  Atomic  Weights  in  Finding  Formulas  and 
Determining  Percentage  Composition.  —  It  is  clear  from 
foregoing  sections  (64-67)  that  atomic  weights  have  a 
fundamental  relation  to  formulas,  molecular  weights,  and 
percentage  composition  of  compounds.  Thus,  we  have 
seen  that  the  composition  of  a  compound  may  be  expressed 
in  per  cent  or  by  a  formula  and  that  the  formula  by 
its  symbols  connects  the  composition  with  the  atomic 
weights.  Mathematically  we  express  composition  in  per 
cent;  chemically  we  express  composition  by  formulas. 
Thus,  the  composition  of  water  may  be  expressed  equally 
well  by  hydrogen  =  11.18  per  cent  and  oxygen  =  88.82 
or  by  H2O. 

(i)  The  calculation  of  a  formula,  when  the  atomic 
weights  and  the  percentage  composition  are  known,  is 
simply  the  process  of  finding  the  small  integral  numbers 
by  which  each  atomic  weight,  as  represented  by  its  symbol, 
must  be  multiplied  in  order  to  express  the  composition. 
The  composition  of  sulphuric  acid  is  hydrogen  =  2.04 
per  cent,  sulphur  =  32.65  per  cent,  oxygen  =  65.31  per 
cent.  If  the  percentage  of  each  element  is  divided  by  the 
corresponding  atomic  weight,  the  quotients  are  2.04,  1.02, 
and  4.08.  Reducing  these  quotients  to  integral  numbers 
(by  dividing  by  1.02),  the  final  quotients  are  2,  1,4.  But 
these  quotients  represent  the  ratio  of  the  atomic  weights 
in  a  molecule;  that  is,  the  relative  number  of  atoms  of 
each  element  in  a  molecule.  And  since  atoms  are  repre- 
sented by  symbols,  the  formula  of  sulphuric  acid  must  be 
H2S04.  The  formula  of  a  compound  calculated  by  this 
method  is  called  its  simplest  formula.  (See  Determination 
of  Molecular  Formulas  of  Compounds,  Chapter  XII.) 


LAW  AND  THEORY  63 

(2)  The  calculation  of  composition  in  per  cent,  or,  as 
it  is  usually  called,  the  percentage  composition  of  a  com- 
pound, is  simply  the  process  of  transposing  the  chemical 
formula  of  a  compound  into  the  equivalent  mathematical 
form.  Let  us  take  an  example.  The  formula  of  potas- 
sium chlorate  is  KC1O3.  This  formula  represents  a  molec- 
ular weight  of  122.5,  i-e-  39  +  35-5  +  4&  =  122.5  (using 
approximate  atomic  weights).  Now  if  the  respective 
parts  of  potassium,  chlorine,  and  oxygen  (viz.  39,  35.5, 
48)  are  divided  by  122.5  and  the  quotient  then  multiplied 
by  100  (e.g.  39/122.5  X  100),  the  product  is  the  per 
cent  of  each  element  in  sulphuric  acid.  It  is  sometimes 
more  convenient  to  solve  the  problem  by  a  proportion. 
Thus,  the  proportions  for  finding  the  percentage  composi- 
tion of  potassium  chlorate  are :  - 

39  :  122.5  ::  x  :  I00>  x  =  3x-^3  Per  cent  °f  potassium 
35.5  :  122.5  ••  x  :  100,  x  =  28.98  per  cent  of  chlorine 
48  :  122.5  ::  x  '•  I00>  x  =  SQ-1^  Per  cent  °f  oxygen 

The   percentage  composition  of  any  compound  can  be 
calculated  from  its  formula  by  this  method. 

EXERCISES 

1.  Define  law  and  theory  as  used  in  science. 

2.  State  the  law  of  the  conservation  of  matter.    Illustrate  it. 

3.  State  the  law  of  constant  composition.    Illustrate  it 

4.  State  the  law  of  multiple  proportions.    Illustrate  it. 

5.  State  the  atomic  theory. 

6.  Discuss  the  relation  of  atoms  to  molecules. 

7.  Describe  a  chemical  change  in  terms  of  the  atomic  theory. 

8.  Interpret  by  the  atomic  theory  the  three  laws:    (a)  conservation  of 
matter,  (6)  constant  composition,  (c)  multiple  proportions. 

9.  What  are  atomic  weights?    Molecular  weights? 

10.  What  is  the   symbol   of   an   element?    Interpret   the   expressions 
H,  20,  N»,  2P,  30,  K2,  S2,  2C1. 

11.  What  is  the  formula  of  a  compound?    What  does  a  formula  repre- 


64  CHEMISTRY 

sent?      Interpret  the  expressions:    H2O,  2H2O,  KNO3,  4H2SO4,  NH4OH, 
C2H5.OH,  3Ca(OH)2,  A12(SO4)3. 

12.  Give  the  symbols  of  the  following  elements:    oxygen,  hydrogen, 
nitrogen,  zinc,  copper,  magnesium,  platinum,  iron,  sodium,  sulphur,  carbon. 

13.  What  elements  correspond  to  the  following  symbols:    Na,  Cu,  K, 
Zn,  S,  P,  Pt,  Pb,  H,  Hg,  Fe,  Mg? 

14.  Give  the  formulas  of  the  following  compounds:    water,  potassium 
chlorate,  sulphuric  acid,  magnesium  oxide,  copper  oxide,  sodium  hydroxide. 

PROBLEMS 

1.  Show  that  the  following  sets  of  compounds  illustrate  the  law  of 
multiple  proportions:    (a)  S  =  50  per  cent  and  O  =  50  per  cent,  S  =  40 
and  O  =  60;     0)  Sn  =  78.8  and  S  =  21.2,   Sn  =  65.02   and  S  =  34.98; 
(c)  Hg  =  84.92  and  Cl  =  15.08,  Hg  =  73.8  and  Cl  =  26.2. 

2.  Calculate  the  molecular  weight  (or  multimolecular  weight)  of  the 
following  compounds  by  finding  the  sum  of  the  atomic  weights:    (a)  mag- 
nesium oxide     (MgO),    (b)  hydrogen    peroxide    (H2O2),    (c)   zinc  chloride 
(ZnCla),     (d)    2Cu(N03)2,     (e)  3A12(SO4)3,     (/)  potassium     ferrocyanide 
(K4Fe(CN)6),  (g)  2Na2B4O7,  (h)  crystallized  ferrous  sulphate  (FeSO4.7H2O). 

3.  Calculate  the  simplest  formula  of  the  compounds  which  have  the 
following  percentage  composition:     (a)  Cl  =  60.68,  Na  =  39.31;    (b)  S  = 
23.52,  Ca  =  29.41,  O  =  47.05;   (c)  C  =  40,  H  =  6.67,  O  =  53.33. 

4.  As  in  Problem  3:    (a)  N  =  26.17,  H  =  7.48,  Cl  =  66.35;    (6)  As  = 
75.8,  O  =  24.2;    (c)  N  =  82.35,  H  =  17.63. 

6.  As  in  Problem  3:  (a)  Si  =  19.5,  C  =  66.62,  H  =  13.88;  (6)  Ca  = 
38.71,  P  =  20,  O  =  41.29;  (c)  H  =  i,  K  =  39.06,  C  =  11.99,  O  =  47.95. 

6.  Calculate  the  formula  of  a  compound  18  gm.  of  which  contain  8.4  gm. 
of  iron  and  9.6  gm.  of  sulphur. 

7.  As  in  Problem  6:    .84  gm.  contain  .587  gm.  of  iron  and  .253  gm.  of 
oxygen. 

8.  Calculate  the  percentage  composition  of  (a)  hydrochloric  acid  (HC1), 
(b)  hydrogen  sulphide  (H2S),  (c)  ammonia  (NH3),  (</)  hydrogen  peroxide. 

9.  As  in  Problem  8:    (a)  Calcium  oxide  (CaO),  (b)  calcium  carbonate 
(CaCO3),  (c)  calcium  sulphate  (CaSO4),  (d)  calcium  fluoride  (CaF2). 

10.  As  in  Problem  8:    (a)  Cane  sugar  (Ci2H22On)  and  (b)  grape  sugar 
(C6H1206). 

11.  As  in  Problem  8:    (a)  Sodium  phosphate  (Na3PO4),  (b)  disodium 
phosphate  (HNa2PO4),  (c}  monosodium  phosphate  (H2NaPO4),  (rf)  phos- 
phoric acid  (H3PO4). 

12. ,  Calculate  the  per  cent  of  (a)  copper  and  (b)  water  in  crystallized 
copper  sulphate  (CuSO4-5H2O). 

13.  Calculate  the  per  cent  of  (a)  F  in  SiF4,  (b)  Al  in  A1PO4,  (c)  O  in 
MnO2,  (d)  Pb  in  PbCO3. 


CHAPTER  VIII     . 

CHEMICAL  REACTIONS,  EQUATIONS,  AND    CALCULATIONS 

69.  Chemical  Reactions.  —  The  substances  that  par- 
ticipate in  chemical  changes  are  said  to  undergo  chemical 
action,  to  interact,  or  to  react.    The  chemical  change  itself 
is  called  a  chemical  reaction,  or  simply  a  reaction. 

Most  of  the  reactions  we  shall  study  can  be  assigned 
to  one  of  four  general  classes  —  decomposition,  combina- 
tion, substitution,  and  double  decomposition.  The  first 
three  have  already  been  defined  and  illustrated  (11, 15, 22). 
The  fourth  class,  double  decomposition,  is  the  decomposi- 
tion of  two  compounds  and  the  subsequent  recombination 
of  their  parts  into  two  other  compounds.  Thus,  when 
sodium  chloride  and  silver  nitrate  are  mixed  in  solution, 
both  compounds  decompose  and  their  parts  recombine  to 
form  silver  chloride  and  sodium  nitrate. 

70.  Chemical  Equations.  —  In  the  chapters  on  Oxy- 
gen, Hydrogen,  and  Water,  certain  chemical  changes  were 
represented  by  equations.    In  these  equations  words  were 
used,  e.g. 

Mercuric  Oxide  =  Mercury  +  Oxygen 

In  the  preceding  chapter  we  found  that  elements  and 
compounds  may  be  represented  by  symbols  and  formulas. 
Therefore,  we  can  now  remodel  the  preliminary  equations 
into  chemical  equations  by  using  symbols  and  formulas 
in  place  of  words.  The  above  equation  then  becomes :  — 

HgO  =  Hg  +  O 


66  CHEMISTRY 

Equations  are  used  so  frequently  in  chemistry  it  is 
necessary  to  study  them  fully.  Their  meaning  can  best 
be  made  clear  by  a  further  discussion  of  one  of  the  four 
classes  of  reactions,  viz.  decomposition.  When  mercuric 
oxide  is  heated,  it  changes  into  mercury  and  oxygen. 
This  reaction  is  expressed  by  the  equation :  — 

HgO  Hg      +      O 

Mercuric  Oxide        Mercury       Oxygen 

This  equation  may  be  read  in  several  ways:  (a)  Mercuric 
oxide  decomposes  into  mercury  and  oxygen;  (b)  one 
molecule  of  mercuric  oxide  by  decomposition  .forms  one 
atom  of  mercury  and  one  atom  of  oxygen;  (c)  216  parts 
by  weight  of  mercuric  oxide  yield  200  parts  of  mercury 
and  1 6  parts  of  oxygen  (since  these  are  the  respective 
molecular  and  atomic  weights) ;  (d)  mercuric  oxide  equals 
mercury  plus  oxygen. 

Let  us  consider  each  lettered  statement.  In  (a)  nothing 
is  said  about  the  physical  conditions  attending  the  chemi- 
cal change  expressed  by  the  equation,  because  ordinary 
chemical  equations  are  not  designed  to  tell  anything  about 
the  physical  conditions  (i.e.  temperature,  physical  state, 
solution,  etc.)  under  which  the  reaction  starts,  proceeds, 
and  ends.  The  statements  in  (b)  and  (c)  are  closely  con- 
nected and  must  be  considered  together.  What  determines 
the  number  of  atoms  and  molecules  in  an  equation?  First, 
equations  are  the  outcome  of  experiments.  Second,  ordi- 
nary chemical  equations  express  quantitative  relations. 
That  is,  they  not  only  emphasize  the  fact  that  a  chemical 
change  conforms  to  the  law  of  the  conservation  of  matter, 
but  they  also  express  the  proportions  by  weight  in  which 
substances  take  part  in  the  reaction.  For  example,  when 
a  given  weight  of  mercuric  oxide  is  decomposed  into  mer- 


REACTIONS,  EQUATIONS,  AND  CALCULATIONS  67 

cury  and  oxygen,  experiment  shows  that  (i)  the  weight 
of  the  mercuric  oxide  equals  the  sum  of  the  weights  of  the 
resulting  mercury  and  oxygen  and  (2)  the  actual  weights 
involved  are  in  the  ratio  of  216  to  200  to  16  respectively. 
These  numbers  necessitate  the  number  of  atoms  and  mole- 
cules used  in  the  equation  HgO  =  Hg  +  O,  because  these 
and  only  these  chemical  expressions  stand  for  216,  200, 
and  1 6  respectively.  Briefly,  ordinary  chemical  equations 
show  by  the  proper  number  of  atoms  and  molecules  the 
result  of  a  reaction.  The  reacting  substances  are  often 
called  the  factors  and  the  final  substances  the  products. 
The  statement  in  (d)  is  sometimes  used  because  chemical 
equations  seem  like  those  used  in  algebra.  But  chemical 
equations  are  not  algebraic;  their  terms  cannot  be  trans- 
posed or  factored  nor  can  the  equations  themselves  be 
predicted.  An  arrow  may  be  used  instead  of  the  equality 

sign,  e.g.  HgO >  Hg  +  O.  Both  the  arrow  and  the 

equality  sign  may  be  read  as  give(s),  form(s),  or  produce(s). 
Sometimes  the  plus  sign  is  read  as  and,  with,  acted  upon 
by,  or  reacting  with.  Each  equation  is  the  outcome  of 
experiment,  and  although  the  equation  contains  algebraic 
signs,  it  has  none  of  the  properties  of  an  algebraic  equation 
except  equality  between  the  total  weights  on  each  side  of 
the  equation. 

Statements  similar  to  (a),  etc.,  can  be  made  about 
equations  corresponding  to  the  other  classes  of  reactions; 
the  number  of  atoms  and  molecules  will  differ,  of  course, 
with  different  reactions  though  it  is  customary  to  use  in 
the  equation  the  least  number  that  will  correctly  express 
the  reaction. 

71.  Writing  Equations.  —  In  order  to  express  a  reac- 
tion by  an  equation,  we  must  know  the  factors  and 
products  of  a  reaction.  For  if  we  know  the  names  of  the 


68  CHEMISTRY 

substances  involved  in  a  reaction,  we  can  (i)  find  their 
symbols  or  formulas  in  the  book,  (2)  construct  a  prelim- 
inary equation,  and  then  (3)  balance  the  equation.  By 
balancing  an  equation  we  mean  selecting  the  proper 
coefficients,  subscripts,  or  both,  so  that  there  shall  be  an 
equal  number  of  atoms  of  each  element  on  both  sides  of 
the  equation.  Some  examples  will  make  this  method 
clear.  When  phosphorus  burns  in  oxygen,  phosphorus 
pentoxide  (see  Index)  is  formed.  The  preliminary  equa- 
tion is :  - 

P  +  O  =  P205 

Here  it  is  evident  that  to  balance  the  equation  we  need 
2P  and  50  on  the  left.  Hence  the  final  equation  is:  - 

2P  +  50  =  P205 

Again,  when  zinc  and  hydrochloric  acid  interact,  hydro- 
gen and  zinc  chloride  are  formed.  The  preliminary  equa- 
tion made  from  the  symbols  and  formulas  is :  - 

Zn  +  HC1  =  H  +  ZnCl2 

By  inspection,  it  is  evident  that  two  atoms  of  chlorine 
are  on  the  right  and  only  one  on  the  left.  To  obtain  the 
C12  it  is  necessary  to  write  2HC1.  But  2HC1  means  not 
only  2C1  but  2H.  Hence  the  equation  becomes  finally :  — 

Zn  +  2HC1  =  2H  +  ZnCl2 

A  final  inspection  shows  that  an  equal  number  of  atoms 
of  each  element  is  on  both  sides  of  the  equation.  Many 
equations  may  be  written  by  applying  this  method  to 
the  facts  found  by  experiment  (see  Exercises  at  the  end  of 
this  chapter).  There  are  other  methods  of  writing  equa- 
tions, but  they  need  not  be  discussed  here. 


REACTIONS,  EQUATIONS,  AND  CALCULATIONS  69 

72.  Equations  for  Preceding  Reactions.  —  The  equations  cor- 
responding to  many  reactions  already  discussed  may  be  collected 
here,  partly  for  review  and  partly  for  future  use. 

HgO         =       Hg  +0 

Mercuric  Oxide       Mercury  Oxygen 

KC103  KC1  +      30 

Potassium  Chlorate      Potassium  Chloride       Oxygen 

PbO2  =       PbO  +  O 

Lead  Dioxide      Lead  Oxide  Oxygen 

BaO2  =        BaO  +  O 

Barium  Dioxide       Barium  Oxide  Oxygen 

Na202  +        H20     =  2NaOH         +         O 

Sodium  Peroxide  Water        Sodium  Hydroxide       Oxygen 

H2O  aH  +  O 

Water        Hydrogen  Oxygen 

S          +        2O      =  SO2 

Sulphur          Oxygen       Sulphur  Dioxide 

C        +        20      =  C02 

Carbon          Oxygen       Carbon  Dioxide 

Cu        +         O      =         CuO 
Copper  Oxygen      Copper  Oxide 

Zn  +          H2SO4  2H         +  ZnSO4 

Zinc  Sulphuric  Acid      Hydrogen  Zinc  Sulphate 

Zn  +  2HC1  2H      -f  ZnCl2 

Zinc  Hydrochloric  Acid      Hydrogen  Zinc  Chloride 

Na  +         H20  =         H  +           NaOH 

Sodium             Water  Hydrogen        Sodium  Hydroxide 

Ca  +       2H2O  2H  +           Ca(OH)2 

Calcium            Water  Hydrogen           Calcium  Hydroxide 

Al          +        3NaOH  =        3H        +  Na3AlO3 

Aluminium       Sodium  Hydroxide        Hydrogen         Sodium  Aluminate 

2H        +       O  H20 

Hydrogen       Oxygen         Water 
CuO          +        2H  H2O       +       Cu 

Copper  Oxide       Hydrogen          Water  Copper 

NaCl  +        AgN03        =          AgCl          +         NaNO3 

Sodium  Chloride       Silver  Nitrate       Silver  Chloride       Sodium  Nitrate 


70  CHEMISTRY 

73.    Chemical    Calculations    based    on    Equations.  - 

Many  problems  may  be  solved  by  reactions.  Obviously, 
any  convenient  weights  of  zinc  and  sulphuric  acid  might 
be  allowed  to  interact,  but  the  factors  and  products  are 
always  in  the  proportions  given  in  the  equation :  - 

Zn  +      H2S04       =       2H      +      ZnSO4 

Zinc        Sulphuric  Acid        Hydrogen        Zinc  Sulphate 

65  98  2  161 

If  zinc  and  sulphuric  acid  are  brought  together  in  any 
other  proportion,  some  of  the  acid  or  the  metal  will  be 
left  over  unused.  This  equation  means  that  zinc  and 
sulphuric  acid  always  interact  in  the  ratio  of  65  to  98, 
and  produce  hydrogen  and  zinc  sulphate  in  the  ratio  of 
2  to  161.  Hence,  if  we  know  the  actual  weight  of  one 
substance  participating  in  a  reaction,  all  other  weights 
involved  can  be  readily  calculated.  Let  us  take  an 
example.  Suppose  45  gm.  of  zinc  interact  with  sul- 
phuric acid;  the  weights  of  (a)  acid  required,  (b)  hydro- 
gen formed,  and  (c)  zinc  sulphate  produced  may  be 
calculated  as  follows :  - 

(1)  Write  the  correct  chemical  equation  for  the  reaction, 
thus:- 

Zn+.H2S04  =  2H  +  ZnS04 

(2)  Place  under  each  term  of  the  equation  its  atomic 
or  molecular  weight,1  as  the  case  may  be,  thus:  - 

Zn  +  H2SO4  =  2  H  +  ZnSO4 
65          98  2          161 

(3)  Place  above   the  proper  terms  the  known  weight 
and  required  weight   (i.e.  x,  y,  z,  etc.)   involved  in  the 
problem,  thus:— 

1  The  atomic  weights  are  given  in  the  table  on  the  inside  of  the  back 
cover.    Molecular  weights  are  obtained  by  adding  the  proper  atomic  weights. 


REACTIONS,  EQUATIONS,  AND  CALCULATIONS     71 

45  x  y  z 

Zn  +  H2S04  =  2  H  +  ZnSO4 
65          98  2  161 

(4)  State  in  the  form  of  a  proportion  the  four  terms 
involved,  remembering  that  the  known  and  required 
weights  are  in  the  same  ratio  as  the  atomic  or  molecular 
weights.  Thus,  the  three  proportions  in  this  problem 
are:  - 

(a)  45  :  x  ::  65  :  98;     x  =  67.8  gm.  sulphuric  acid. 

(b)  45  :  y  ::  65  :  2;       y  =  1.38  gm.  hydrogen. 

(c)  45  :  z  ::  65  :  161;    z  =  111.4  gm.  zinc  sulphate. 

EXERCISES 

1.  What  is  a  chemical  equation?    What  are  the  factors  and  products  in 
an  equation?    Illustrate  your  answer. 

2.  Select  an  equation  from  72  and  read  it  in  different  ways. 

3.  Select  from  72  equations  illustrating  the  four  classes  of  reactions. 

4.  Interpret    the   equation   Mg  -f  O  =  MgO   by   stating    (a)  what   it 
means  and  (6)  what  it  does  not  include  or  express. 

6.   What  do  the  plus  (+)  and  equality  (  =  )  signs  mean  in  equations? 

6.  Write  equations  for  the  following  reactions   (see   71):    (a)  Calcium 
and  hydrochloric  acid  form   calcium  chloride  and  hydrogen,     (b)  Potas- 
sium sulphate  and  barium  chloride  form  barium  sulphate  and  potassium 
chloride,     (c)  Calcium  carbonate  and  hydrochloric  acid  form  calcium  chlo- 
ride, water,  and  carbon  dioxide. 

7.  As  in  Exercise  6:   (a)  Calcium  oxide  and  carbon  dioxide  form  cal- 
cium carbonate,     (b)  Chlorine  and  phpsphorus  form  phosphorus  trichloride. 
(c)  Carbon  and  lead  oxide  (PbO)  form  lead  and  carbon  monoxide. 

8.  State   all   that  the  following  equations  mean:    (a)   BaCl2  +  H2SO4 
=  BaS04  +  2  HC1;    (b)  Pb(NO3)2  +  H2S  =  PbS  +  2HNO3;    (c)  A1CU  + 
3NH4OH  =  A1(OH),  +  3NH4C1;  (d)  Ca(OH)2  +  CO2  =  CaCO3  +  H2O. 


72  CHEMISTRY 

PROBLEMS 

1.  How  many  grams  of  oxygen  can  be  prepared  from  (a)  45  gm.  of 
mercuric  oxide,  (b)  i  kg.  of  potassium  chlorate,  (c)   1000  gm.  of  water? 

2.  As  in  Problem   i,  from  (a)  750  gm.  of  lead  dioxide  (PbO2),  (6)  a 
metric  ton  of  barium  dioxide  (BaO2),  (c)  37  gm.  of  sodium  peroxide  (by 
interaction  with  water)? 

3.  What  volume  of  oxygen  at  standard  conditions  could  be  obtained 
from  10  gm.  of  potassium  chlorate? 

4.  How  many  grams  of  potassium  chlorate  (92  per  cent  pure)  are  needed 
to  prepare  (a)  100  gm.  of  oxygen  and  (b)   100  1.  (at  standard  conditions)? 

5.  Suppose  85   gm.  of  water  are  decomposed,    (a)  What  weights  and 
(b)  what  volumes  of  gases  are  produced? 

6.  If  a  metric  ton  of  pure  carbon  is  burned  in  air,  what  weights  of  other 
substances  are  involved? 

7.  One  gram  of  copper  is  heated  intensely  in  air,  and  the  product  is 
reduced  by  a  gas.    Calculate  the  weights  of  the  other  substances  involved 
in  the  two  reactions. 

8.  Hydrogen  is  prepared  from  sulphuric  acid  and  40  gm.  of  zinc.    Cal- 
culate the  weights  of  the  products  of  the  reaction. 

9.  Calculate  the  required  weights  involved  in  the  following  reactions: 

(a)  water  and  100  milligrams  of  sodium,  (b)  calcium  and  100  milligrams  of 
water,  (c)  sodium  hydroxide  and  25  gm.  of  aluminium. 

10.  What  weight  of  carbon  dioxide  is  formed  by  burning  a  metric  ton 
of  coal  which  is  90  per  cent  carbon? 

11.  If  a  balloon  holds   150  kg.  of  hydrogen,  how  much  (a)  zinc  and 

(b)  sulphuric  acid  are  needed  to  generate  the  gas? 

12.  Sixty  grams  of  mercuric  oxide  are  decomposed.     What  volume  of 
oxygen  at  91°  C.  and  380  mm.  is  produced? 

13.  How  much  water  is  in  (a)  34  gm.  of  crystallized  zinc  sulphate 
(ZnSO4.7H2O),  (b)  1000  kg.  of  selenite  (CaSO4.2H2O),  (c)  1000  gm.  washing 
soda  crystals  (Na2CO3.ioH2O)? 

14.  The  interaction  of  barium  nitrate  and  sodium  sulphate  is  expressed 
by    (BaNO3)2  +  Na2SO4  =  BaSO4  +  2NaNO3.      If    170    gm.    of    barium 
nitrate  are  used,  calculate  the  weights  of  the  other  compounds  involved. 

16.  Ammonia  gas  and  hydrogen  chloride  form  solid  ammonium  chloride. 
Write  the  equation  for  this  reaction.  If  210  gm.  of  ammonia  are  used, 
calculate  the  weights  of  the  other  compounds  involved. 

16.  The  oxygen  is  liberated  from  10  gm.  of  potassium  chlorate,  and  10 
gm.  of  sulphur  are  burned  in  the  gas.  How  much  sulphur,  if  any,  is  left? 


CHAPTER   IX 


CHLORINE  — HYDROCHLORIC  ACID  — ACIDS,  SALTS, 
AND  BASES 

Chlorine 

74.  Occurrence.  —  Free    chlorine   is   never   found   in 
nature,  but  its  compounds  are  widely  distributed,   the 
most  abundant  being  sodium  chloride  or  common  salt. 
Many  compounds  of  chlorine  with  potassium,  magnesium, 
and  calcium  are  found  in  the  deposits  at  Stassfurt  in 
Germany  (377).     The  salts  found  in  sea  water   contain 
about  55  per  cent  of  chlorine. 

75.  Preparation.  —  Chlorine  is  prepared  on  a  large  scale 
by  the  electrolysis  of  a  solution  of 

sodium  chloride.  When  a  current  of 
electricity  is  passed  through  a  solu- 
tion of  sodium  chloride,  chlorine  gas 
is  liberated  in  one  compartment  of 
the  apparatus  and  sodium  hydroxide 
is  formed  in  the  other.  The  chlorine 
is  conducted  off  through  pipes,  and 
the  dissolved  sodium  hydroxide  is  Fis-  22-  ~  Apparatus 

rr      .   .    _,  ,          rr*  .  for  the  Preparation  of 

drawn  off  at  intervals.     This  process        chlorine  by  the  Elec 
is  further  described  in  371.  trolysis  of  a  Solution 

of  Sodium  Chloride. 
This  process  may  be  readily  demonstrated. 

The  apparatus  is  shown  in  Fig.  22.  A  solution  of  sodium  chloride 
is  put  in  the  battery  jar  A;  a  little  litmus  solution  is  added  and 
then  enough  dilute  hydrochloric  acid  to  color  the  solution  a  distinct 
red.  A  wooden  block  (B)  divides  the  jar  into  two  compartments 


74  CHEMISTRY 

(C  and  D),  and  the  two  pieces  of  electric  light  carbon  serve  as 
electrodes  (E  and  F).  Soon  after  the  current  (from  four  or  more 
cells  or  from  a  reduced  street  current)  is  turned  on,  the  solution  is 
bleached  by  the  liberated  chlorine  in  one  compartment  and  turned 
blue  by  the  sodium  hydroxide  in  the  other.  The  chlorine  can  also 
be  detected  by  its  odor. 

Chlorine  is  prepared  in  the  laboratory  by  heating  a 
mixture  of  manganese  dioxide  and  concentrated  hydro- 
chloric acid.  (See  Part  II,  Exp.  33.)  The  equation  is:  - 

MnO2    +       4HC1      =     2C1    +    MnCl2     +  2H2O 

Manganese        Hydrochloric        Chlorine        Manganese          Water 
Dioxide  Acid  Bichloride 

76.  Physical  Properties.  —  Chlorine  is  a  greenish 
yellow  gas.  Its  color  suggested  the  name  chlorine  (from 
the  Greek  word  chloros,  meaning  greenish  yellow),  which 
was  given  to  it  by  Davy  about  1810.  It  has  a  disagreeable, 
suffocating  odor,  which  is  very  penetrating.  If  breathed, 
it  irritates  the  sensitive  lining  of  the  nose  and  throat, 
and  a  large  quantity  would  cause  death.  It  is  about 
2.5  times  heavier  than  air.  Hence  it  is  easily  collected  by 
downward  displacement,  i.e.  by  conducting  it  downward 
to  the  bottom  of  a  vessel  and  allowing  it  to  displace  the 
air.  A  liter  of  dry  chlorine  at  o°  C.  and  760  mm.  weighs 
3.22  gm.  Chlorine  can  be  readily  liquefied  and  solidified. 

Water  dissolves  chlorine.  The  solution  is  yellowish,  and 
smells  strongly  of  chlo  ine.  Chlorine  water,  as  the  solu- 
tion is  called,  is  unstable.  If  the  solution  is  placed  in 
the  sunlight,  oxygen  is  soon  liberated  and  hydrochloric 
acid  is  formed  (53,  Fig.  18);  intermediate  changes  occur, 
but  the  simplest  equation  for  the  essential  change  is:  - 

H20  +     2C1     =      2HC1      +     O 

Water        Chlorine        Hydrochloric        Oxygen 
Acid 


CHLORINE  75 

77.  Chemical  Properties.  —  Many  elements  unite  vig- 
orously with  chlorine.  Thus,  if  sodium,  iron  (thread), 
copper  (wire),  or  other  metals  are  heated  and  then  put 
into  chlorine,  they  burn;  the  sodium  produces  a  dazzling 
light,  and  the  copper  and  iron  glow  and  emit  dense  fumes. 
These  chemical  changes  illustrate  the  broad  use  of  the 
term  combustion.  (Compare  17,  24.)  The  compound 
formed  in  each  case  is  a  chloride,  i.e.  a  compound  of  chlo- 
rine and  one  other  element.  Chlorine  combines  readily 
with  hydrogen.  Hence,  a  jet  of  burning  hydrogen  when 
lowered  into  chlorine  continues  to  burn,  forming  a  color- 
less gas  called  hydrogen  chloride,  which  appears  as  a 
white  cloud,  especially  when  the  breath  is  blown  gently 
across  the  mouth  of  the  vessel.  The  simplest  equation 
for  this  reaction  is :  — 

H       +      Cl      =  HC1 

Hydrogen         Chlorine         Hydrogen  Chloride 

The  tendency  of  chlorine  to  combine  with  hydrogen  is 
so  great  that  the  hydrogen  of  many  compounds  is  with- 
drawn chemically  by  chlorine.  Chlorine  does  not  com- 
bine with  carbon;  hence  substances  which  contain  carbon 
burn  in  chlorine  with  a  smoky  flame. 

Chlorine  bleaches.  This  property  depends  upon  the 
fact  that  chlorine  interacts  with  water  and  ultimately 
liberates  free  oxygen;  the  latter  then  decomposes  the 
complex  coloring  matter  into  colorless  substances.  If 
an  envelope  on  which  the  postmark  or  a  pencil  mark  is 
still  visible  is  placed  in  moist  chlorine,  these  marks  will 
not  be  bleached  because  they  are  largely  carbon;  but 
the  writing  ink,  which  is  substantially  a  compound  of 
hydrogen,  carbon,  and  iron,  will  disappear.  Litmus 
paper  and  many  kinds  of  colored  cloth  are  bleached  by 


76  CHEMISTRY 

moist  chlorine.  Dry  chlorine  does  not  bleach.  (See  Part 
II,  Exp.  33.) 

78.  Bleaching  Powder  is  the  main  source  of  the  chlorine 
used  in  the  bleaching  industries.     It  is  sometimes  called 
chloride  of  lime.    It  is  a  yellowish  white  substance  which 
smells  like  chlorine.     Acids,   like   sulphuric  and   hydro- 
chloric acid,  liberate  from  bleaching  powder  its  "  available 
chlorine,"  which  is  about  37  per  cent.    The  equations  for 
the  interaction  of  acids  and  bleaching  powder  may  be 
written  thus :  - 

CaOCl2  +  H2S04    =     2C1     +  CaSO4  +  H2O 

Bleaching        Sulphuric        Chlorine        Calcium        Water 
Powder  Acid  Sulphate 

CaOCl2  +     2HC1      =  2C1  +  CaCl2  +  H2O 

Hydrochloric  Calcium 

Acid  Chloride 

Bleaching  powder  is  manufactured  by  passing  chlorine 
gas  over  dry  calcium  hydroxide,  the  equation  for  the 
reaction  being:  — 

Ca(OH)2       +    2C1    =       CaOCl2  -f      H2O 

Calcium  Hydroxide       Chlorine       Bleaching  Powder       Water 

79.  Bleaching.  —  Immense     quantities     of     bleaching 
powder  are  used  to  whiten  cotton  and  linen  goods  and 
paper  pulp.     The  pieces  of  yellowish,  unbleached  cloth 
are  drawn  by  machinery  through    numerous  vats  con- 
taining weak  solutions  of  bleaching  powder  and  of  acid, 
and  finally  washed  to  remove  all   traces  of  substances 
which  would  injure  the  fabric. 

Bleaching  is  chemically  an  oxidizing  process.  The  oxy- 
gen when  it  is  liberated  from  water  by  chlorine  is  said 
to  be  in  the  nascent  state.  This  means  that  the  gas  is 
exceedingly  active,  because  it  is  not  only  uncombined,  but 


CHLORINE  77 

just  ready  to  unite  with  those  elements  for  which  it  has  a 
marked  tendency  to  combine.  Hence  this  nascent  oxygen 
readily  decomposes  the  colored  substances  and  changes 
them  into  colorless  compounds.  (See  Part  II,  Exp.  34.) 

80.  Uses  of  Chlorine.  —  Besides  the  use  of  chlorine  in 
the  manufacture  of  bleaching  powder,  large  quantities  of 
the  gas  are  made  into  useful  compounds  of  chlorine,  e.g. 
carbon  tetrachloride   (CC14)   which  is  used  in  "  pyrene  " 
fire  extinguishers  and  as  a  solvent  for  grease;   the  non- 
inflammable  cleaning  mixture  called  "  carbona "  contains 
carbon  tetrachloride. 

81.  Chlorides  are  formed  when  chlorine  combines  with 
other  elements,  just  as  oxides  are  formed  from  oxygen. 
(See  also  89.)     Two  equations  are:- 

Na     +      Cl  NaCl 

Sodium          Chlorine  Sodium  Chloride 

Sb       +     3C1  SbCla 

Antimony  Antimony  Trichloride 

Hydrochloric  Acid 

82.  Hydrochloric  Acid  is  the  common  name  of  a  water 
solution  of  hydrogen  chloride,  HC1.    Hydrogen  chloride 
is  a  gas,  which  is  very  soluble  in  water.    This  solution  is 
known  in  commerce  as  muriatic  acid  (from  the  Latin  word 
muria,  meaning  brine),  but  it  is  more  properly  called 
hydrochloric   acid.      Hydrogen   chloride   is   often   called 
hydrochloric  acid  gas. 

83.  Preparation.  —  The  gas   is  prepared   by  heating 
sulphuric  acid  and  sodium  chloride.     If  the  mixture  is 
gently  heated,  the  chemical  change  is  represented  thus:  — 

NaCl  +  H2S04   =     HC1     +  HNaSO4 

Sodium        Sulphuric        Hydrogen        Acid  Sodium 
Chloride  Acid  Chloride  Sulphate 


78  CHEMISTRY 

But  at  a  high  temperature  the  equation  for  the  reaction 


2NaCl  +  H2SO4  =  2HC1 


Na2SO4 

Sodium  Sulphate 


The  solution  is  prepared  by  passing  the  gas  into  water. 
(See  Part  II,  Exp.  35.) 

84.    Commercial  Hydrochloric  Acid  is  manufactured  in  enormous 
quantities  by  the  method  used  in  the  laboratory  (Fig.  23).      The 

mixture  of  salt  and  sul- 
phuric acid  is  put  into 
the  cast  iron  retort  A 
and  heated  by  the  fur- 
nace B  to  a  moderate 
temperature;  as  soon 
as  the  mass  becomes 
pasty  it  is  raked  out 
upon  the  flat  heater  A' 
and  heated  to  a  high 
temperature  by  the  fur- 


Fig.  23.  —  Apparatus  for  the  Manufacture 
of  Hydrochloric  Acid. 


nace  B  .  The  hydrogen  chloride  escapes  through  C  and  C'  into 
an  absorbing  tower  rilled  with  resistent  material  over  which  water 
trickles;  as  the  gas  passes  up  the  tower,  it  is  absorbed  by  the  de- 
scending water,  and  the  solution  flows  out  at  the  bottom  as  con- 
centrated acid.  Hydrochloric  acid  is  produced  in  England  as  a 
by-product  in  the  manufacture  of  sodium  carbonate  by  the  Leblanc 
process  (365). 

85.  Physical  Properties.  —  Hydrogen  chloride  is  a 
colorless  gas.  It  has  a  choking,  sharp  taste,  and  irritates 
the  lining  of  the  nose  and  throat.  The  gas  does  not  burn 
nor  support  combustion.  It  is  about  1.25  times  heavier 
than  air.  A  liter  at  o°  C.  and  760  mm.  weighs  1.64  gm. 
The  gas  becomes  a  colorless  liquid  when  subjected  to 
pressure  and  a  moderately  low  temperature.  The  extreme 
solubility  of  hydrogen  chloride  in  water  is  one  of  its  most 


CHLORINE  79 

striking  properties.  When  it  escapes  into  moist  air,  it 
forms  \vhiu-  fumes  which  are  really  minute  drops  of  a 
solution  of  the  gas  in  the  moisture  of  the  air.  At  ordinary 
temperatures  about  500  1.  of  gas  dissolve  in  i  1.  of  water. 
The  solution  is  the  familiar  hydrochloric  acid;  its  specific 
gravity  is  about  1.2,  and  it  contains  approximately  40  per 
cent  (by  weight)  of  hydrogen  chloride.  The  gas  readily 
escapes,  hence  the  acid  forms  fumes  when  exposed  to  air. 
86.  Chemical  Properties.  —  Perfectly  dry  hydrogen 
chloride  has  little  or  no  chemical  activity.  The  moist 
gas  unites  readily  with  certain  substances,  e.g.  ammonia 
gas;  in  this  case  dense  white  clouds  of  ammonium  chloride 
are  formed,  the  equation  for  the  reaction  being:  — 

HC1    +    NHa          NH4C1 

Hydrogen       Ammonia       Ammonium 
Chloride  Chloride 

This  reaction  is  sometimes  used  as  a  test  for  hydrogen 
chloride  (or  for  ammonia) .  Hydrochloric  acid  —  the 
water  solution  of  hydrogen  chloride  —  has  marked  chemi- 
cal properties.  Like  most  members  of  the  important  class 
of  compounds  called  acids,  it  has  a  sour  taste  and  reddens 
blue  litmus;  it  also  reacts  with  many  metals,  liberating 
hydrogen  and  forming  chlorides  of  the  metals,  thus:- 

Zn    +    2HC1      =      2H    +    ZnCl2 

Zinc       Hydrochloric       Hydrogen         Zinc 

Acid  Chloride 

It  also  forms  chlorides  by  interaction  with  oxides  and 
hydroxides  of  metals,  thus:- 

CaO     +     2HC1  =  CaCl2  +  H2O 

Calcium  Oxide  Calcium       Water 

Chloride 

NaOH     +     HC1  =        NaCl     +     H2O 

Sodium  Hydroxide  Sodium  Chloride       Water 


8o  CHEMISTRY 

87.  Aqua  Regia.  —  Hydrochloric  acid  and  nitric  acid 
interact  and  liberate  chlorine,  thus:  — 

3HC1    +    HNO3  =     2C1     +   NOC1  +  2H2O 

Hydrochloric          Nitric          Chlorine          Nitrosyl          Water 
Acid  Acid  Chloride 

A  mixture  of  one  volume  of  concentrated  nitric  acid  and 
three  volumes  of  concentrated  hydrochloric  acid  is  usually 
used.  If  such  a  mixture  is  added  to  a  metal,  a  chloride 
of  the  metal  is  formed.  The  alchemists  gave  it  the 
name  aqua  regia,  meaning  "  royal  water,"  to  emphasize 
the  fact  that  it  dissolves  the  "noble"  metal  gold.  (See 
Part  II,  Exp.  43.) 

88.  Volumetric    Composition    of    Hydrogen    Chloride.  —  Experi- 
ments show  that  hydrogen  chloride  is  composed  of  hydrogen  and 
chlorine  in  the  ratio  of  i  :  i  by  volume.     When  a  mixture  of  equal 
volumes  of  hydrogen  and  chlorine  is  exposed  to  the  direct  sunlight 
or  to  the  action  of  an  electric  spark,  the  gases  combine,  hydrogen 
chloride  is  formed  with  no  residue,  and  the  volume  of  the  result- 
ing gas  equals  the  sum  of  the  volumes  of  hydrogen  and  chlorine 
used.     The  volumetric  relations  of  hydrogen,  chlorine,  and  hydro- 
gen chloride  may  be  expressed  by:  — 

i  volume  of  hydrogen  +  i   volume  of  chlorine  =  2  volumes  of 
hydrogen  chloride. 

89.  Chlorides  are  compounds  of  chlorine  and  other 
elements.     They  are  formed,  as  we  have  already  seen, 
by  the  direct  combination  of  chlorine  and  metals  and  by 
the  interaction  of  hydrochloric  acid  with  metallic  oxides 
or  hydroxides  (81).     Most  chlorides  are  soluble  in  water. 
The  chlorides  of  lead  and  silver  (and  one  of  the  chlorides 
of  mercury)  are  not;   they  are  formed  as  insoluble  solids 
when  hydrochloric  acid  or  a  soluble  chloride  undergoes 
double  decomposition  with  a  soluble  lead  or  silver  com- 
pound.    Thus:  — 


CHLORINE  8 1 

Pb(N03)2  +      2HC1      =      PbCl2     +  2HN03 

Lead  Nitrate        Hydrochloric        Lead  Chloride         Nitric 
Acid  Acid 

AgNO3    +  NaCl  =  AgCl  +  NaNO3 

Silver  Sodium  Silver          Sodium 

Nitrate  Chloride        Chloride       Nitrate 

The  formation  of  insoluble  solids  by  double  decompo- 
sition (and  certain  other  changes)  is  called  precipitation, 
and  the  solid  itself  is  called  a  precipitate.  Precipitates 
often  have  properties  which  are  readily  determined.  Thus, 
silver  chloride  is  white  and  curdy,  and  soon  turns  purple 
in  the  light;  moreover  it  dissolves  in  ammonium  hydroxide 
owing  to  the  formation  of  a  complex  soluble  compound, 
which,  however,  is  transformed  by  dilute  nitric  acid  into 
silver  chloride.  Other  chlorides  have  different  properties. 
Hence,  the  precipitation  of  silver  chloride  serves  as  a  test 
for  hydrochloric  acid  and  soluble  chlorides. 

A  molecule  of  a  chloride  may  contain  one  or  more 
atoms  of  chlorine,  and  occasionally  the  name  of  the  com- 
pound indicates  this  fact.  E.g.  manganese  dichloride 
(MnCl2),  antimony  trichloride  (SbCls).  If  a  metal  forms 
two  chlorides,  the  two  are  distinguished  by  modifying 
the  name  of  the  metal;  the  one  containing  the  smaller 
proportion  of  chlorine  ends  in  -ous,  that  containing  the 
larger  in  -ic.  Thus,  mercurous  chloride  is  HgCl,  but 
mercuric  chloride  is  HgCl2. 

Acids,  Salts,  and  Bases 

90.  Acids.  —  Hydrochloric  acid,  as  already  stated,  is 
a  member  of  an  important  class  of  compounds  called 
acids.  The  properties  of  this  acid  are  characteristic  of 
the  class.  All  acids  contain  hydrogen  that  can  be  re- 


82  CHEMISTRY 

placed  by  certain  metals;  and  the  compound  formed  by 
this  replacement  is  called  a  salt.  Other  compounds 
contain  hydrogen,  but  they  are  not  classed  as  acids  unless 
they  form  salts  by  replacement  of  the  hydrogen  by  a 
metal.  Thus,  water  and  sugar  contain  hydrogen,  but  the 
hydrogen  in  water  does  not  form  a  salt  by  replacement  of 
its  hydrogen  by  a  metal,  nor  can  the  hydrogen  in  sugar 
be  replaced  by  a  metal.  Furthermore,  most  acids  have 
a  sour  taste  and  redden  blue  litmus.  Substances  which 
act  thus  on  blue  litmus  are  said  to  have  an  acid  reaction. 
The  presence  of  acids  is  often  conveniently  detected  by 
these  simple  tests.  For  example,  vinegar,  pickles,  many 
fruits,  and  some  wines  have  a  sour  taste  and  turn  blue 
litmus  red;  further  examination  reveals  the  acids  in  these 
substances.  (See  Part  II,  Exps.  37,  44.) 

91.  Salts,  as  a  rule,  are  not  sour  and  do  not  redden 
blue  litmus.  Substances  which  act  thus  on  litmus  are 
often  described  as  having  a  neutral  reaction.  Many 
salts  have  the  taste  associated  with  a  familiar  member  of 
this  class,  viz.  common  salt  or  sodium  chloride.  This 
class  of  compounds  has  many  members  and  their  prop- 
erties are  somewhat  varied.  Salts  invariably  contain  a 
metal  and  a  non-metal,  and  most  salts  also  contain  oxy- 
gen. Chlorides  are  salts.  Thus,  we  have  referred  to 
the  chlorides  of  zinc,  sodium,  calcium,  lead,  silver. 
Reference  has  also  been  made  to  some  salts  containing 
oxygen,  e.g.  potassium  chlorate,  KC1O3,  and  sodium 
sulphate,  Na2SO4;  these  compounds  are  salts  of  chloric 
acid  and  sulphuric  acid  respectively.  In  discussing  the 
chemical  properties  of  hydrochloric  acid,  it  was  stated  that 
chlorides  are  formed  from  hydrochloric  acid  not  only  by 
the  interaction  with  metals  but  also  with  oxides  and  hy- 
droxides of  metals  (89).  That  is,  the  metal,  whatever  its 


CHLORINE  83 

source,  replaces  the  hydrogen  of  the  acid  and  thereby 
forms  a  salt.  The  formation  of  salts  from  acids  and  hy- 
droxides is  very  important.  (See  Part  II,  Exp.  39.) 

92.  Bases.  —  Hydroxides  belong  to  another  important 
class  of  compounds  called  bases.     Solutions  of  most  bases 
turn  red  litmus  blue  —  just  the  opposite  of  acids,  and  are 
said  to  have  a  basic  or  an  alkaline  reaction.    Bases  always 
contain  a  metal  united  with  oxygen  and  hydrogen,  e.g. 
sodium  hydroxide,  NaOH.      (See  Part  II,  Exps.  38,  44.) 

93.  Neutralization.  —  Now  acids,  salts,  and  bases  have 
fundamental  chemical  relations.     When  we  mix  solutions 
containing  weights  of  an  acid  and  a  base  proportional  to 
their  molecular  weights,  the  acid  and  base  interact  com- 
pletely;   the  final  solution  has  none  of  the  characteristic 
properties  of  an  acid  or  a  base,  but  it  does  have  the  prop- 
erties of  a  salt.     That  is,  the  acid  and  base  destroy  the 
marked  properties  of  each  other  and  form  a  salt.     The 
acid  and  base  neutralize  each  other.     For  example,  when 
hydrochloric  acid  and  sodium  hydroxide  interact,  sodium 
chloride  and  water  are  formed.     The  equation  for  the 
reaction  is :  — 

HC1  +  NaOH  =  NaCl  +  H2O 

Hydrochloric    Sodium         Sodium        Water 
Acid         Hydroxide       Chloride 

A  chemical  change  in  which  an  acid  and  a  base  neutralize 
each  other  and  form  a  salt  and  water  is  called  neutraliza- 
tion. In  neutralization  the  hydrogen  and  oxygen  of  the 
base  act  as  a  unit.  This  group  of  atoms  (OH)  is  called 
hydroxyl.  Hydroxyl  does  not  exist  free  and  uncombined 
like  elements  and  compounds,  but  it  acts  like  a  single 
atom  in  many  chemical  changes.  It  is  called  a  radical. 
(See  Part  II,  Exp.  40.) 
Neutralization  illustrates  double  decomposition.  In 


84  CHEMISTRY 

the  chemical  change  just  cited  both  the  hydrochloric 
acid  and  the  sodium  hydroxide  are  decomposed  and  their 
parts  are  recombined  in  a  different  way;  i.e.  sodium 
chloride  and  water  are  the  new  compounds  resulting 
from  the  recombination  (69). 

For  the  present,  we  may  regard  salts  as  compounds 
formed  either  from  acids  and  bases  by  neutralization,  or 
from  acids  by  the  substitution  of  a  metal,  directly  or 
indirectly,  for  the  hydrogen  of  the  acid.  The  nature  and 
interrelation  of  acids,  salts,  and  bases  will  be  further 
discussed  (see  Chapter  XIV). 

94.  Naming  Acids  and  Salts.  —  Three  other  acids 
besides  hydrochloric  acid  contain  no  oxygen,  viz.  hydro- 
fluoric (HF),  hydrobromic  (HBr),  and  hydriodic  (HI). 
The  corresponding  salts  end  in  ide,  e.g.  chloride,  fluoride, 
bromide,  and  iodide.  Sometimes  the  compound  com- 
monly known  as  hydrogen  sulphide  (H2S)  is  called  an 
acid,  and  its  salts  are  the  sulphides  (274).  Oxygen  is  a 
constituent  of  most  acids  and  salts,  and  the  names  of 
the  oxy-acids  and  oxy-salts  are  related,  especially  the 
suffixes.  This  relation  can  be  best  illustrated  by  the 
chlorine  acids  that  contain  oxygen.  These  acids  are 
hypochlorous  (HC1O),  chlorous  (HC1O2),  chloric  (HC1O3), 
and  perchloric  (HC1O4).  In  forming  the  names  of  the 
corresponding  salts  the  suffix  ous  becomes  tie,  while  ic 
becomes  ate;  the  prefixes  are  not  changed.  Thus,  the 
names  of  the  corresponding  sodium  salts  are  sodium 
hypochlorite,  chlorite,  chlorate,  and  perchlorate  respect- 
ively. In  the  case  of  the  acids  of  some  elements  the 
body  of  the  name  is  modified,  e.g.  sulphuric  becomes 
sulphate,  phosphoric  phosphate,  and  tartaric  tartrate. 

Bases  are  distinguished  by  placing  the  name  of  the  metal  before 
the  word  hydroxide,  e.g.  sodium  hydroxide  (NaOH),  calcium  hydroxide 
(Ca(OH)2). 


CHLORINE  85 

EXERCISES 

1.  What  useful  compounds  contain  chlorine? 

2.  Sketch  from  memory  the  apparatus  used  to  prepare  chlorine. 

3.  Summarize  the  physical  properties  of  chlorine.     How  can  chlorine  be 
quickly  distinguished  from  the  gases  previously  studied? 

4.  Summarize  the  chemical  properties  of  chlorine. 

6.   Develop  the  topics:    (a)  nascent  state,  (b)  chlorine  water,  (c)  liquid 
chlorine,  (d)  chlorine  is  an  oxidizing  agent. 

6.  What  is  (a)  hydrogen  chloride,  (b)  muriatic  acid,  (c)  chloride  of  lime, 
(d)  hydrochloric  acid,  (e)  commercial  hydrochloric  acid? 

7.  Give  the  name  and  formula  of  five  chlorides. 

8.  State  the  characteristic  properties  of  hydrogen  chloride. 

9.  Summarize  the  chief  properties  of  hydrochloric  acid. 

10.  What  is  the  volumetric  composition  of  hydrogen  chloride? 


PROBLEMS 

1.  Calculate  the  weight  of  chlorine  in   (a)  2  kg.  of  sodium  chloride, 
(b)  2  mg.  of  calcium  chloride,  (c)  i  metric  ton  of  aluminium  chloride,  and 
(d)  45  gm.  of  potassium  chlorate. 

2.  (a)  What  is  the  weight  of  15  1.  of  chlorine  gas  measured  at  20°  C.  and 
790  mm.?     (6)  How  many  grams  of  potassium  chloride  are  needed  to  pre- 
pare the  weight  of  chlorine  found  in  (a)? 

3.  How  many  grams  of  each  of  the  three  products  are  formed  when 
hydrochloric  acid  interacts  with  85  grams  of  manganese  dioxide? 

4.  How  many  liters  of  (a)  chlorine  and  (b)  hydrogen  chloride  can  be 
obtained  from  27  gm.  of  sodium  chloride?     (Standard  conditions.) 

6.  How  many  cubic  centimeters  of  hydrochloric  acid  solution  having  a 
specific  gravity  of  1.21  and  containing  42.06  per  cent  of  HC1  by  weight,  are 
needed  to  form  4  liters  of  chlorine?  (Standard  conditions.) 

6.  (a)  How  much  sodium  chloride  can  be  formed  by  burning  sodium  in 
40  gm.  of  chlorine?     In  40  liters?     (b)  How  much  antimony  trichloride  can 
be  formed  from  the  same  quantities  of  chlorine?     (Standard  conditions.) 

7.  (a)  What  is  the  weight  of  25  1.  of  hydrogen  chloride  at  18°  C.  and 
765  mm.?     (b)  What  volume  at  the  same  temperature  and  pressure  will 
25  gm.  occupy? 

8.  How  much  hydrogen  chloride  by  weight  and  by  volume  (at  standard 
conditions)  can  be  obtained  from  a  metric  ton  of  sodium  chloride  which  is 
85  per  cent  pure? 


CHAPTER  X 


NITROGEN  — AMMONIA  — NITRIC  ACID  AND  NITRATES - 
NITROGEN  OXIDES 

Nitrogen 

96.  Occurrence.  —  Nitrogen  is  an  essential  ingredient 
of  the  atmosphere,  being  about  four  fifths  (or  exactly 
78.122  per  cent)  of  this  vast  mixture  of  gases  which 
envelops  the  earth.  Besides  being  a  constituent  of  nitric 
acid,  nitrates,  and  ammonia,  and  of  the  many  compounds 
related  to  them,  nitrogen  is  also  found  in  many  animal  and 
vegetable  substances  fundamentally  related  to  life,  e.g.  the 
compounds  called  proteins  (see  Protein,  Chapter  XVII). 
96.  Preparation.  —  Nitrogen  is  prepared  on  a  large 
scale  from  liquid  air  (134).  It  can  also 
be  obtained  from  air  by  removing  the 
oxygen  by  phosphorus  (Fig.  24). 

Phosphorus  is  put  in  a  small  dish  or  a  cru- 
cible cover  supported  on  a  cork  floating  in  a 
vessel  of  water.      Upon  igniting  the  phospho- 
rus with  a  hot  wire  and  placing  a  bell  jar  over 
the  cork,  the   phosphorus  and   oxygen  unite, 
forming  clouds  of  white  phosphorus  pentoxide 
Fig.    24.  —  Prepara-    (P2O6);   this  solid  soon  dissolves  in  the  water, 
tion  of  Nitrogen  by    which  rises  inside  the  jar  owing  to  the  removal 
burning    Phospho-    of  the  oxygen,  and  the  nitrogen  is  finally  left, 
rus  in  Confined  Air. 

It  is  prepared  in  the  laboratory  by 
heating  a  solution  of  sodium  nitrite  and  ammonium  chlo- 
ride. The  equations  expressing  the  reactions  are  - 


NITROGEN  87 

NaNO2  +   NH4C1    =  NH4NO2  +  NaCl 

Sodium          Ammonium        Ammonium        Sodium 
Nitrite  Chloride  Nitrite  Chloride 

NH4N02  =      2N     +  2H20 

Ammonium        Nitrogen         Water 
Nitrite 

Small  quantities  of  nitrogen  are  easily  obtained  by  heating 
ammonium  dichromate  ((NH4)2Cr207).  (See  Part  II, 
Exps.  45,  52.) 

97.  Physical  and  Chemical  Properties  of  Nitrogen.  - 
Nitrogen  is  a  colorless  gas,  and  has  no  taste  or  odor.  It 
is  a  little  lighter  than  oxygen  and  air.  A  liter  at  stan- 
dard conditions  weighs  1.25  gm.  It  is  only  slightly 
soluble  in  water.  Nitrogen  does  not  support  combustion 
nor  sustain  life.  Flames  are  extinguished  by  nitrogen 
and  animals  die  in  it,  because  the  supply  of  oxygen  is  cut 
off.  Subjected  to  a  low  temperature  and  increased  pres- 
sure, nitrogen  becomes  a  colorless  liquid  and  ultimately 
a  white  solid.  (See  Part  II,  Exp.  45  B.) 

The  fact  that  nitrogen  quickly  extinguishes  a  candle  flame  and 
kills  a  mouse  was  first  observed  by  Rutherford,  a  Scottish  physi- 
cian, who  discovered  the  gas  in  1772.  Soon  after,  Lavoisier 
showed  the  true  relation  of  nitrogen  to  the  atmosphere. 

Nitrogen  is  much  less  active  chemically  than  oxygen; 
it  is  sometimes  called  an  inert  element,  because  it  does 
not  combine  with  elements  at  ordinary  temperatures. 
At  high  temperatures  and  under  special  conditions,  how- 
ever, nitrogen  forms  many  compounds.  It  combines  with 
magnesium  and  a  few  other  metals  at  a  red  heat,  forming 
nitrides,  e.g.  magnesium  nitride  (Mg3N2).  Electric 
sparks  cause  nitrogen  to  combine  with  oxygen  and  with 
hydrogen,  forming  nitric  oxide  (NO)  and  ammonia 


88 


CHEMISTRY 


98.  Relation  of  Nitrogen  to  Life.  —  Nitrogen,  as  well  as 
oxygen,  is  vitally  connected  with  life,  though  in  a  different 
way.  All  animals  need  nitrogen  for  their  growth.  Now 
although  we  live  in  an  atmosphere  containing  such  a  large 
proportion  of  this  gas,  we  cannot  assimilate  it  directly. 
The  nitrogen  needed  by  animals  must  be  in  chemical  com- 
bination to  become  available.  That  is,  it  must  be  eaten 
in  the  form  of  nitrogenous  food,  such  as  lean  meat,  fish, 
wheat  and  other  grains  (see  Protein,  Chapter  XVII). 
Nor  have  plants,  with  few  exceptions,  power  to  assimi- 
late free  nitrogen  from  the  atmosphere.  Most  plants  take 
up  combined  nitrogen  from  the  soil  in  the  form  of  nitrates 
or  of  ammonia.  Hence  combined  nitrogen  is  being  con- 
stantly removed  from  the 
soil,  and  in  order  to  restore 
it,  some  nitrogen  compound 
must  be  added,  e.g.  sodium 
nitrate  (NaNOs),  calcium 
nitrate  (Ca(NO3)2),  ammo- 
nium chloride  (NH4C1),  or 
ammonium  sulphate 
((NH4)2SO4) ;  organic  ma- 
terials are  often  used,  e.  g. 
manure,  dried  blood,  and 
meat  or  fish  scraps.  Such  a 
replenishing  substance,  or  a 
mixture  containing  it,  is 
called  a  fertilizer.  Many  experiments  have  shown,  how- 
ever, that  leguminous  plants,  such  as  peas,  beans,  and 
clover,  take  up  nitrogen  from  the  air  by  means  of  bacteria, 
which  are  in  nodules  on  their  roots  (Fig.  25).  This  process 
is  called  fixation  of  nitrogen.  Sometimes  soil  is  treated  with 
a  preparation  which  contains  nitrogen-forming  bacteria. 


Fig.  25.  —  A  Leguminous  Plant 


AMMONIA  89 

Ammonia 

The  term  ammonia  includes  both  the  gas  (NH3)  and  its 
solution  in  water  (NH4OH),  though  the  latter  is  more 
accurately  called  ammonium  hydroxide. 

99.  Formation  of  Ammonia.  —  When   vegetable    and 
animal  matter  containing  nitrogen  decays,  the  nitrogen 
and  hydrogen  are  usually  liberated  as  ammonia.    The  odor 
of  ammonia  may  be  noticed  near  stables.     If  animal  sub- 
stances containing  nitrogen  are  heated  (especially  with 
lime  or  soda-lime),  ammonia  is  given  off.     Soft  coal  con- 
tains combined  nitrogen  and  hydrogen,  and  when  the  coal 
is  heated  to  make  illuminating  gas,  ammonia  is  obtained. 

100.  Preparation.  —  Ammonia  gas  is  prepared  in  the 
laboratory  by  heating  ammonium  chloride  with  a  base, 
usually  calcium  hydroxide.    The  equation  for  the  reac- 
tion is  — 

2NH4C1  +  Ca(OH)2  =  2NH4OH  +  CaCl2 

Ammonium  Calcium  Ammonium         Calcium 

Chloride  Hydroxide  Hydroxide         Chloride 

The  ammonium  hydroxide  is  unstable,  especially  when 
heated,  and  decomposes  into  ammonia  and  water,  thus:  — 

NH4OH  =     NH3     +  H20 

Ammonium  Hydroxide         Ammonia        Water 

The  gas  is  very  volatile,  and  is  usually  collected  by  up- 
ward displacement,  i.e.  by  allowing  the  gas  to  flow  upward 
into  a  bottle  and  displace  the  air.  The  solution  is  pre- 
pared by  conducting  the  gas  into  water.  (See  Part  II, 
Exps.  46,  53.) 

The  ammonia  gas  from  which  the  ammonia  of  commerce  is 
manufactured  is  obtained  mainly  from  the  ammoniacal  liquor  or 
gas  liquor  of  the  illuminating  gas  works.  The  gases  which  come 


QO  CHEMISTRY 

from  the  retorts  in  which  the  coal  is  heated  are  washed  with  water. 
This  impure  gas  liquor  is  treated  with  lime  to  liberate  the  am- 
monia, which  is  absorbed  in  tanks  containing  hydrochloric  acid  or 
sulphuric  acid.  This  solution  upon  the  addition  of  a  base  (e.g. 
calcium  hydroxide)  gives  up  its  ammonia,  which  is  dissolved  in 
distilled  water,  forming  thereby  the  ammonium  hydroxide  or  am- 
monia of  commerce.  Some  ammonia  is  obtained  from  the  gases 
liberated  from  coke  ovens,  and  it  is  also  manufactured  by  the 
direct  combination  of  nitrogen  and  hydrogen. 

101.  Physical  Properties  of  Ammonia.  —  Ammonia 
gas  is  colorless.  It  has  an  exceedingly  pungent  odor,  and 
if  inhaled  suddenly  or  in  large  quantities,  it  brings  tears  to 
the  eyes  and  may  cause  suffocation.  It  is  a  light,  volatile 
gas,  being  only  .59  times  as  heavy  as  air.  A  liter  of  the 
gas  at  o°  C.  and  760  mm.  weighs  .77  gm.  Ammonia  gas 
is  easily  liquefied  if  cooled  to  o°  C.  and  subjected  to  a 
pressure  of  4.2  atmospheres  (121).  Liquefied  ammonia 
is  often  called  anhydrous  ammonia,  because  it  contains  no 
water.  It  boils  at  about  —34°  C.  Hence,  if  it  is  exposed 
to  the  air  or  warmed  in  any  way,  it  changes  into  the  gas, 
and  in  so  doing  absorbs  considerable  heat.  This  fact  has 
led  to  the  extensive  use  of  liquid  ammonia  in  the  manu- 
facture of  ice  (104). 

Ammonia  gas  is  very  soluble  m- water.  A  liter  of  water 
at  o°  C.  dissolves  1148  1.  of  gas  (measured  at  o°  C.  and 
760  mm.),  while  at  the  ordinary  temperature  i  1.  of  water 
dissolves  about  700  1.  of  gas.  This  solution  of  the  gas 
is  often  called  ammonia,  though  other  names,  e.g.  am- 
monium hydroxide  and  ammonia  water  are  sometimes 
applied  to  it  (105);  it  gives  off  the  gas  freely,  when 
heated,  as  may  easily  be  discovered  by  the  odor  or  by  the 
formation  of  dense  white  fumes  of  ammonium  chloride 
(NH4C1)  when  the  solution  is  exposed  to  hydrochloric 
acid.  The  commercial  solution  called  ammonia  13  lighter 


AMMONIA  91 

than  water  (its  specific  gravity  being  about  .88)  and  con- 
tains approximately  35  per  cent  (by  weight)  of  the  com- 
pound NH3.  (See  Part  II,  Exp.  46.) 

102.  Chemical  Properties  of  Ammonia.  —  Ammonia 
gas  will  not  burn  in  air  under  ordinary  conditions,  nor  will 
it  support  combustion  as  the  term  is  usually  used  ;  but  if 
the  air  is  heated  or  if  its  proportion  of  oxygen  is  increased, 
a  jet  of  ammonia  gas  will  burn  in  it  with  a  yellowish  flame. 
When  electric  sparks  are  passed  through  ammonia  gas, 
it  is  decomposed  into  nitrogen  and  hydrogen;  the  reac- 
tion is  incomplete,  however,  for  nitrogen  and  hydrogen 
unite  to  some  extent  under  the  same  conditions.  Am- 
monia reacts  with  certain  elements.  Dried  ammonia  gas 
and  heated  magnesium  form  hydrogen  and  magnesium 
nitride,  thus:  - 


2NH3    +      3Mg      =         Mg3N2     +      6H 

Ammonia         Magnesium         Magnesium  Nitride 

Ammonia  and  chlorine  interact,  thus:  — 

NH3     +     3C1     =      N      +         3HC1 

Ammonia         Chlorine         Nitrogen         Hydrochloric  Acid 

Ammonia  combines  directly  with  water,  forming  am- 
monium hydroxide,  thus:  - 

NH3    +  H20  =          NH4OH 

Ammonia        Water        Ammonium  Hydroxide 

It  also  combines  with  certain  gases,  e.g.  hydrogen  chloride 
(HC1),  thereby  forming  ammonium  chloride  (NH4C1). 
This  reaction  serves  as  a  test  for  ammonia  gas. 

103.  Composition  of  Ammonia  Gas.  —  Experiments 
show  that  ammonia  is  a  compound  of  nitrogen  and  hydro- 
gen, especially  its  synthesis  from  these  elements.  By 
utilizing  the  fact  that  ammonia  and  chlorine  react  and 


92  CHEMISTRY 

liberate  nitrogen,  the  volumetric  composition  of  am- 
monia can  be  shown  to  be  nitrogen  is  to  hydrogen  as  i  13. 
A  supplementary  experiment  shows  that  two  volumes 
of  ammonia  are  formed  by  the  union  of  one  volume  of 
nitrogen  and  three  volumes  of  hydrogen.  (Compare 
54,  88.) 

In  demonstrating  the  volumetric  composition  of  ammonia,  a  tube 
A  (Fig.  26)  filled  with  a  known  volume  of  chlorine  is  provided  with 
a  funnel  B  through  which  concentrated  ammonium 
BQ  hydroxide  is  slowly  dropped  into  the  chlorine,  until 

the  reaction  ceases.      After  the  excess  of  ammonium 
A  a  hydroxide  is  neutralized  with  sulphuric  acid,  the  vol- 

ume of  nitrogen  left  is  found  to  be  one  third  of  the 
original  volume  of  chlorine.  Now  hydrogen  and 
chlorine  combine  in  equal  volumes  (88).  Hence  the 
volume  of  hydrogen  withdrawn  from  the  ammonia 
must  be  equal  to  the  original  volume  of  chlorine. 
But  this  volume  is  three  times  the  volume  of  the 
nitrogen,  therefore  there  must  be  three  times  as  much 
hydrogen  as  nitrogen  in  ammonia  gas. 

104.   Ammonia  as  a  Refrigerant.  —The  use 
of  ammonia  in  producing  low  temperatures 

Fig.     26  —    depends  upon  the  fact  that  liquefied  ammo- 
Apparatus       j  r  . \  •  ? 
for   deter-    ma  (n°t  ordinary  ammonia)  changes  rapidly 

mining  the    into  a  gas  when  its  temperature  is  raised  or 
Composi-    the  pressure  reduced.    Hence,  if  liquefied  am- 

tionofAm-  ,,  ,          a  .  .         . 

moniaGas  moma  1S  allowed  to  now  through  a  pipe  im- 
mersed in  a  solution  of  sodium  chloride  or 
calcium  chloride  (technically  called  a  brine),  the  ammonia 
evaporates  in  the  pipe  and  cools  the  brine,  which  may  be 
used  directly  as  a  refrigerant  or  for  making  ice.  In  some 
cold-storage  plants,  breweries,  packing  houses,  and  sugar 
refineries,  this  cold  brine  is  pumped  through  pipes  placed 
in  the  rooms  where  a  low  temperature  is  desired. 


AMMONIA 


93 


The  construction  and  general  operation  of  an  ice-making  plant 
is  shown  in  Fig.  27.      Liquefied  ammonia  is  forced  from  a  tank 

Cold  Water  trickling  over  the 

ammonia  pipes  to  condense 

the  compressed  sras 


Expansion  valve 


Brine  pump 


Fig.  27.  —  Apparatus  for  utilizing  Cooled  Brine  as  a  Refrigerant  and  in 
making  Ice. 

into  a  series  of  pipes  which  are  submerged  in  a  large  vat  nearly 
filled  with  brine.  Metal  cans  containing  pure  water  to  be  frozen 
are  immersed  in  the  brine,  which  is  kept  below  the  freezing  point 
of  water  by  rapid  evaporation  of  the  ammonia  in  the  pipes.  After 
several  hours  the  water  in  the  cans  is  frozen  into  cakes  of  ice. 
Sometimes  the  brine  is  circulated  through  pipes,  as  in  a  cold  stor- 
age plant.  As  fast  as  the  ammonia  gas  forms  in  the  pipes,  it  is 
removed  by  exhaust  pumps  into  another  tank,  where  it  is  con- 
densed into  liquefied  ammonia  and  conducted,  as  needed,  into  the 
first  tank  ready  for  renewed  use. 

105.  .Ammonium  Hydroxide  and  Ammonium  Com- 
pounds. —  When  ammonia  gas  is  passed  into  water,  the  am- 
monia combines  with  the  water  to  some  extent  and  forms 
a  solution  of  an  unstable  compound  having  the  composi- 
tion represented  by  the  formula  NH^OH.  This  compound 
is  ammonium  hydroxide.  Ammonium  hydroxide  is  a  base 
(92).  Like  other  members  of  this  class  of  substances  it 
turns  litmus  blue;  it  also  neutralizes  acids,  thus:  - 

NH4OH  +       HC1       =    NH4C1    +  H2O 

Ammonium         Hydrochloric         Ammonium         Water 
Hydroxide  Acid  Chloride 


94  CHEMISTRY 

Ammonium  hydroxide  is  widely  used  as  a  cleansing 
agent  (especially  for  the  removal  of  grease),  as  a  re- 
storative in  cases  of  fainting  or  of  inhaling  irritating 
gases;  large  quantities  are  consumed  in  dyeing  and  calico 
printing,  and  in  the  manufacture  of  dyestuffs,  sodium 
carbonate  and  bicarbonate,  and  ammonium  compounds. 
Its  salts  have  many  domestic,  industrial,  and  agricultural 
uses. 

It  is  believed  that  ammonium  compounds  contain  a 
group  of  atoms  which  acts  chemically  like  an  atom  of  a 
metal.  This  group  is  called  ammonium,  and  its  formula 
is  NH4.  Ammonium  has  never  been  separated  from  its 
compounds.  Ammonium  is  called  a  radical,  because  it  is 
the  root  or  foundation  of  a  series  of  compounds.  It  is 
likewise  called  a  hypothetical  metal,  because  its  exist- 
ence is  assumed,  and  it  acts  chemically  like  metals. 

Nitric  Acid 

106.  Formation. —  Nitric  acid,  HNO3,  is  formed  in  small 
quantities  when  electric  sparks  are  passed  through  moist 
air.    Hence  nitric  acid  or  its  salts  can  be  detected  in  the 
atmosphere  after  a  thunderstorm.'    This  chemical  change 
is  now  being  applied  on  a  commercial  scale  in  Norway. 
Air  is  forced  through  a  large  tube  containing  a  powerful 
electric  arc  spread  out  into  a  disk.    The  nitrogen  gas  and 
oxygen  combine  and  the  gases  are  absorbed  in  water  or  in 
a  solution  of  lime,  thereby  forming  nitric  acid  or  calcium 
nitrate.     The  latter  is  used  as  a  fertilizer  in  place  of 
sodium  nitrate  (98,  373). 

107.  Preparation.  —  Nitric    acid    is   prepared    in    the 
laboratory  and  on  a  large  scale  by  heating  concentrated 
sulphuric   acid  with   a  nitrate,   usually   sodium  nitrate. 


NITRIC    ACID 


95 


Fig.  28  shows  the  appa- 
ratus used  in  the  labo- 
ratory. About  equal 
weights  of  sodium  ni- 
trate and  concentrated 
sulphuric  acid  are  put 
into  the  glass  retort  and 
gently  heated;  the  nitric 
acid  distils  into  the  re- 
ceiver, which  is  kept  cool 
by  water,  ice,  or  moist 
paper.  The  chemical 
change  at  a  low  tem- 
perature is  expressed  by  the  equation  - 


Fig.  28.  —  Apparatus  for  preparing 
Nitric  Acid  in  the  Laboratory. 


NaN03      +      H2SO4      =    HNO3    +         HNaSO4 

Sodium  Nitrate       Sulphuric  Acid       Nitric  Acid       Acid  Sodium  Sulphate 

But  if  the  temperature  is  high  and  an  excess  of  the  nitrate 
is  present,  the  equation  is:  - 


2NaN03 


H2S04    =    2HNO3 


Na*SO 


At  a  high  temperature  part  of  the  nitric  acid  decomposes; 
hence  excessive  heat  is  usually  avoided.  (See  Part  II, 
Exp.  47.) 

Nitric  acid  is  manufactured  in  an  iron  retort  connected  with  glass 
or  earthenware  tubes  in  which  the  vapor  is  condensed  by  running 
water;  the  tubes  are  inclined  and  so  arranged  that  the  nitric  acid 
runs  back  into  a  reservoir  from  which  it  may  be  drawn  as  needed 
(Fig.  29). 

108.  Physical  Properties.  —  Pure  nitric  acid  is  a  color- 
less liquid,  but  the  commercial  acid  is  yellow  or  reddish, 
due  to  absorbed  nitrogen  compounds,  chlorine,  or  iron 


96 


CHEMISTRY 


compounds.  The  acid  that  has  been  exposed  to  the 
sunlight  is  often  brown,  and  if  the  light  is  intense,  a 
brownish  gas  may  often  be  seen  in  bottles  of  the  acid. 
It  is  somewhat  volatile,  and  the  vapor  dissolves  readily  in 


Fig.   29.  —  Apparatus  for  the  Manufacture  of  Nitric  Acid. 

water;  hence  the  acid  forms  irritating  fumes  when  exposed 
to  air,  especially  moist  air.  (Compare  85.)  The  specific 
gravity  of  the  commercial  acid  is  about  1.4,  and  it  con- 
tains about  65  per  cent  of  the  compound  HNO3,  the  rest 
being  water. 

109.  Chemical  Properties.  —  A  solution  of  nitric  acid 
has  the  properties  characteristic  of  that  class  of  com- 
pounds. It  is  sour,  turns'  blue  litmus  red,  and  forms 
salts  —  the  nitrates.  It  is  an  unstable  substance,  and 
decomposes  readily;  among  the  decomposition  products 
is  a  brown  gas,  nitrogen  dioxide  (NO2),  which  causes  the 
brown  color  referred  to  above  (108). 

Nitric  acid  reacts  readily  with  many  substances.  With 
organic  compounds  like  the  skin  it  forms  a  yellow  com- 
pound. Hence  it  stains  the  skin  yellow,  and  the  concen- 
trated acid  causes  serious  burns.  With  other  organic 
compounds  it  forms  explosives,  such  as  nitroglycerin  and 
nitrocellulose  (230).  One  of  the  decomposition  products 


NITRIC  ACID  97 

of  nitric  acid  is  oxygen.  Usually  the  oxygen  is  not  liber- 
ated but  oxidizes  whatever  is  available.  Hence  nitric 
acid  is  an  oxidizing  agent.  Thus,  charcoal  burns  bril- 
liantly in  the  hot  acid,  while  straw,  sawdust,  hair,  and 
similar  substances  are  charred  and  even  inflamed  by  it; 
some  organic  compounds,  when  heated  with  nitric  acid, 
are  completely  decomposed  into  carbon  dioxide  and  water. 
In  the  mixture  of  nitric  and  hydrochloric  acids  called 
aqua  regia,  nitric  acid  acts  as  an  oxidizing  agent  (87). 
Finally,  it  interacts  readily  and  often  violently  with 
metals,  metallic  oxides  and  hydroxides.  (Compare  86, 
89.)  The  products  of  these  reactions  vary,  the  essential 
ones  being  nitrates  and  nitrogen  oxides  (112).  (See  Part 
II,  Exps.  48,  54.) 

110.  Uses.  —  Nitric  acid  is  one  of  the  common  labora- 
tory acids.    Large  quantities  are  used  in  the  manufacture 
of  nitrates,  dyestuffs,  sulphuric  acid,  and  explosives,  and 
in  etching  copper  plates. 

111.  Composition  of  Nitric  Acid.  —  Many  independent   experi- 
ments show  that  the  composition  of  nitric  acid  is  expressed  by  the 
formula  HNO3.      (i)   When  electric  sparks  are  passed  through  a 
bottle  containing  moist   air  or  a  solution  of  potassium  hydroxide, 
the  water  becomes  acid  to  litmus  or  the  liquid  will  be  found  to 
contain  a  trace  of  potassium  nitrate.      (2)  Nitric  acid  may  be.  re- 
duced to  ammonia  by  nascent  hydrogen,  thus  showing  that  the 
acid  contains  nitrogen.      (3)  Conversely,  if  a  mixture  of  ammonia 
and  air  is  passed  over  a  mass  of  hot,  porous  platinum,  nitric  acid 
is  formed.      (4)  If   the  acid  is  allowed  to  flow  through  a  hot  por- 
celain or  clay  tube,  oxygen  is  one  of  the  gaseous  products.      (5) 
Analysis  shows  that  it  contains   1.59  per  cent  of  hydrogen,  22.22 
of  nitrogen,  and  76.19  of  oxygen. 

112.  Nitrates.  —  Nitric  acid  forms  salts  called  nitrates. 
They  are  prepared  by  the  methods  usually  used  for  salts, 
i.e.  the  interaction  of  nitric  acid  and  metals  or  metallic 


98  CHEMISTRY 

oxides '  and  the  neutralization  of  hydroxides  by  nitric 
acid.  Nitrates  are  formed  in  the  soil  by  the  slow  action 
of  bacteria  on  complex  nitrogen  compounds;  this  process, 
which  is  called  nitrification,  is  slow  but  very  important, 
since  most  plants  take  up  nitrogen  in  the  form  of  nitrates 
and  utilize  the  nitrogen  in  forming  protein.  The  inter- 
action of  nitric  acid  and  most  metals  is  exceedingly  vig- 
orous, and  for  this  reason,  probably,  the  alchemists  called 
the  acid  aqua  Jortis  —  strong  water.  The  reaction  varies 
with  the  metal,  the  concentration  of  the  acid,  and  the 
temperature.  Hydrogen  is  never  liberated  so  that  it  can 
be  collected,  for  it  is  oxidized  by  the  nitric  acid. 

The  interaction  between  nitric  acid  and  two  different 
metals  will  serve  as  examples  of  the  common  reactions. 
When  moderately  dilute  nitric  acid  is  poured  upon  cop- 
per, a  reddish  brown  gas  is  given  off,  and  the  liquid  turns 
blue,  owing  to  dissolved  copper  nitrate.  The  complete 
equation  for  the  reaction  is :  - 

3Cu    +   8HN03  =  3Cu(N03)2  +  2NO  +  4H2O 

Copper  Nitric  Copper  Nitric 

Acid  Nitrate  Oxide 

This  equation  is  really  made  up  of  three  equations  and 
conceals  the  way  in  which  the*  reactions  really  take 
place.  The  first  reaction  may  be  expressed  thus:  - 

(i)  2HN03  =  30      +      2NO      +  H20 

Nitric  Oxide 

The  oxygen  next  oxidizes  the  copper,  thus:  - 
(2)  3Cu  +  30  =      3CuO 

Copper  Oxide 

The  copper  oxide  then  reacts  with  the  nitric  acid,  thus:  - 
(3)  3CuO  +  6HN03.  =  *  3Cu(N03)2  +  3H2O 

Copper  Nitrate 


NITRIC  ACID  99 

Since  36  is  formed  in  (i)  and  used  in  (2),  and  3CuO  like- 
wise in  (2)  and  (3),  these  two  terms  should  not  appear  in 
the  complete  equation;  the  other  terms  make  up  the 
complete  equation.  In  the  case  of  zinc,  the  reactions  are 
represented  thus:  - 

(1)  3Zn  +  6HN03  =  3Zn(N03)2  +  °H 

Zinc  Nitrate 

(2)  6H  +  2HN03  =  2NO  +  4H20 

The  complete  equation,  from  which  the  common  factor 
(6H)  is  eliminated,  is:  - 

3Zn  +  8HN03  =  3Zn(N03)2  +      2NO      •+  4H2O 

Zinc       Nitric  Acid        Zinc  Nitrate        Nitric  Oxide         Water 

Nitric  oxide  is  represented  as  a  product  of  the  interaction 
of  nitric  acid  and  the  two  metals  copper  and  zinc.  If 
the  reaction  takes  place  in  an  open  vessel,  the  nitric 
oxide,  which  is  a  colorless  gas,  combines  with  oxygen 
and  forms  the  reddish  brown  nitrogen  dioxide  gas.  The 
equation  is:  - 

NO       +     O     =          N02 

Nitric  Oxide        Oxygen        Nitrogen  Dioxide 

Hence  we  often  speak  of  nitrogen  dioxide  as  a  product  of 
the  interaction  of  nitric  acid  and  metals,  though  it  is 
usually  a  secondary  product. 

Most  nitrates  are  white  solids;  those  of  copper,  nickel, 
and  cobalt  are  blue,  green,  and  dark  red  respectively. 
Their  solutions  are  frequently  used  in  the  laboratory. 
The  solids  behave  in  various  ways  when  heated.  Equa- 
tions illustrating  typical  reactions  are :  - 

NaN03  =  NaN02  +  O;  Cu(NO3)2  =  CuO  +  2NO  +  3O; 
NH4N03    =    N20   +   2H20 

Nitrous  Oxide 


ioo  CHEMISTRY 

Since  many  nitrates,  when  heated,  give  up  oxygen,  they 
are  powerful  oxidizing  agents.  Thus,  when  potassium 
nitrate  is  dropped  on  hot  charcoal,  the  charcoal  burns 
vigorously  and  rapidly.  This  kind  of  chemical  action  is 
called  deflagration.  (See  Part  II,  Exps.  50,  51,  54.) 

113.  The  Test  for  Nitrates  (and  of  course  for  nitric 

acid)  is  as  follows:  Add  to  the  nitric  acid  or 
the  solution  of  the  nitrate  an  equal  volume 
of  concentrated  sulphuric  acid,  and  cool  the 
mixture.    Upon  the  cool  mixture  pour  care- 
fully a  cold,  dilute  solution  of  freshly  prepared 
ferrous  sulphate.    A  dark  brown  layer  appears 
Fig.  30.— The    where  the  two  liquids  meet,  owing  to  the  for- 
Test  for  Ni-    matiOn  of  a  brown  unstable  compound  whose 

tricAcidand  .   .  »        *  i  '  :•     i?  o^       ^r\ 

Nitrates.         composition  is  approximately  3FeSO4.2NO. 
(Fig.  30.)     (See  Part  II,  Exp.  49.) 

114.  Nitrous  Acid,  HN02,  is  not   easily   obtained  in   the  free 
state  owing   to    its  instability,   but   its   salts  —  the   nitrites  —  are 
well    known.       Potassium     nitrite    (KNO2)    and    sodium    nitrite 
(NaNO2)  are  formed  by  removing  the  oxygen  from  the  correspond- 
ing nitrate  by  heating  alone  or  with  lead.      Nitrites  give  off  brown 
fumes  (NO2)  when  treated  with  sulphuric  acid,  and  are  thus  readily 
distinguished  from  nitrates.      (See  Part  II,  Exp.  55.) 

Oxides  of  Nitrogen 

115.  Nitrous  Oxide,  N2O,  is  prepared  by  heating  am- 
monium nitrate.     The  equation  for  the  reaction   is:  - 

NH4N03  N20        +        2H20 

Ammonium  Nitrate         Nitrous  Oxide  Water 

This  colorless  gas  has  a  faint  but  pleasant  odor.  It  is  less 
soluble  in  hot  than  in  cold  water,  and  the  solution  has  a 
sweet  taste.  It  is  easily  liquefied  by  reducing  the  tem- 
perature and  applying  pressure,  and  is  often  used  in  this 


OXIDES  OF  NITROGEN  101 

form  to  furnish  the  gas  itself.  The  gas  does  not  burn, 
but  it  supports  the  combustion  of  many  well  burning  sub- 
stances, though  not  so  vigorously  as  oxygen  does.  Thus, 
sulphur,  unless  well  ignited,  will  not  burn  in  nitrous  oxide. 
The  most  striking  property  of  nitrous  oxide  is  its  effect 
on  the  human  system.  If  inhaled  for  a  short  time,  it 
causes  more  or  less  nervous  excitement,  often  manifested 
by  laughter,  and  on  this  account  the  gas  was  called 
"laughing  gas"  by  Davy,  who  first  studied  its  properties 
in  1799.  If  breathed  in  large  quantities,  it  produces 
unconsciousness  and  insensibility  to  pain.  The  gas  is 
often  used  as  an  anesthetic  in  dentistry.  (See  Part  II, 
Exp.  56.) 

116.  The  Volumetric   Composition  of   Nitrous  Oxide.  —  Experi- 
ment shows  that  two  volumes  of  nitrogen  and  one  of  oxygen  form 
two  volumes  of  nitrous  oxide. 

117.  Nitric  Oxide,  NO,  is  usually  prepared  by  the  inter- 
action of  copper  and  dilute  nitric  acid  (sp.  gr.  1.2).    The 
complete  equation  for  the  reaction  is  — 


3Cu    +   8HNO3   =        2NO      +  3 

Copper        Nitric  Acid        Nitric  Oxide        Copper  Nitrate 

Nitric  oxide  is  a  colorless  gas.  It  is  a  little  heavier  than 
air  and  only  slightly  soluble  in  water.  Upon  exposure  to 
air,  it  combines  at  once  with  the  oxygen,  forming  reddish 
brown  fumes  of  nitrogen  dioxide.  The  simplest  equation 
for  this  reaction  is  - 

NO      +     O      =      NO2 

Nitric  Oxide  Nitrogen  Dioxide 

This  property  distinguishes  nitric  oxide  from  all  other 
gases.  It  does  not  burn  nor  support  combustion  unless 
the  burning  substance  (e.g.  phosphorus  or  sodium) 


102  CHEMISTRY 

introduced  is  hot  enough  to  decompose  the  gas  into 
nitrogen  and  oxygen,  and  then,  of  course,  the  liberated 
oxygen  assists  the  combustion.  (See  Part  II,  Exp.  50.) 

118.  The  Volumetric   Composition  of  Nitric  Oxide.  —  By  experi- 
ment it  is  found  that  one  volume  of  nitrogen  and  one  volume  of 
oxygen  form  two  volumes  of  nitric  oxide. 

119.  Nitrogen  Dioxide,  NO2,  is  the  reddish  brown  gas 
formed  by  the  direct  combination  of  nitric  oxide  and 
oxygen.     Thus:  - 

NO      +      O     =        NO2 

Nitric  Oxide  Nitrogen  Dioxide 

It  is  also  produced  by  heating  certain  nitrates.    Thus  — 
Pb(N03)2   =   2N02  +  PbO  +     O 

Lead  Nitrate        Nitrogen         Lead         Oxygen 
Dioxide          Oxide 

The  fumes  of  nitrogen  dioxide  usually  appear  when  nitric 
acid  and  metals  interact,  but,  as  already  stated,  the  nitro- 
gen dioxide  is  produced  by  a  second  reaction,  viz.  the 
combination  of  nitric  oxide  with  the  oxygen  of  the  air. 
Nitrogen  dioxide  has  a  disagreeable  odor,  and  if  breathed 
in  moderately  large  quantities,  it  is  poisonous.  It  inter- 
acts with  water  and  yields  under  ordinary  conditions 
nitric  oxide  and  nitric  acid,  thus :  - 

3N02  +  H20  =  2HN03  +  NO 

Nitrogen        Water  Nitric          Nitric 

Dioxide  Acid  Oxide 

(See  106.)  It  also  dissolves  in  concentrated  nitric  acid, 
forming  fuming  nitric  acid,  which  is  a  powerful  oxidizing 
agent.  (See  Part  II,  Exp.  50.) 

When  the  reddish  brown  gas  is  cooled,  it  gradually  loses  its 
color  and  at  about  26°  C.  becomes  a  yellow  gas,  which  has  the 
composition  represented  by  the  formula  N2O4  and  is  called  nitrogen 


OXIDES  OF  NITROGEN  103 

tetroxide.  Upon  heating  nitrogen  tetroxide,  the  brown  gas  reap- 
pears, and  at  about  140°  C.  the  gas  is  nitrogen  dioxide.  Above 
140°  C.  the  brown  color  fades  owing  to  the  decomposition  of  nitro- 
gen dioxide  into  nitric  oxide  and  oxygen.  A  simple  demonstration 
of  the  relation  between  nitrogen  dioxide  and  tetroxide  is  readily 
made  by  collecting  some  nitrogen  dioxide  in  a  glass  tube,  closing 
the  tube,  and  immersing  the  lower  half  in  ice  water.  The  gas  in 
the  lower  part  becomes  yellow-brown,  whereas  in  the  upper  part  it 
remains  reddish  brown. 

EXERCISES 

1.  Discuss  nitrogen  as  to  (a)  preparation  and  (b)  properties.    Compare 
the  chemical  properties  of  nitrogen  and  oxygen. 

2.  Give  several  tests  for  ammonia. 

3.  How  is  ammonia  gas  liquefied?     What  is  (a)  liquid  ammonia,  (b) 
anhydrous  ammonia,  (c)  liquefied  ammonia?   Describe  the  manufacture  of 
ice  by  liquid  ammonia. 

4.  Starting  with  soft  coal,  state  how  ammonium  sulphate  can  be  manu- 
factured. 

6.   What  is  the  volumetric  composition   of   ammonia  gas? 

6.  State  the  test  for  nitric  acid  and  nitrates. 

7.  Essay  topics:    (a)  A  cold  storage  plant.     (6)  The  cycle  of  nitrogen, 
(c)  Assimilation  of  nitrogen  by  plants,     (d)  Uses  of  nitric  acid. 

PROBLEMS 

1.  What  is  the  weight  of  70  1.  of  nitrogen  at  760  mm.  and  50°  C.? 

2.  What  volume  would  20.4  cc.  of  nitrogen  (measured  over  water)  at 
21.2°  C.  and  763.1  mm.  occupy  at  o°  C.  and  760  mm.? 

3.  What  weight  and  what  volume  (at  o°  C.  and   760  mm.)  of  nitro- 
gen can  be  obtained  from  25  gm.  of  ammonium  nitrite?     (Equation  is 
NH4N02  =  2N  +  2H20.) 

4.  What  is  the  weight  of  32  1.  of  ammonia  gas  measured  at  20°  C.  and 
763  mm.?    What  volume  will  32  gm.  of  ammonia  gas  occupy  at  the  same 
temperature  and  pressure? 

6.   To  what  weight  and  what  volume  of  NH3  are  25  gm.  of  ammonium 
chloride  equivalent?     (Standard  conditions.) 

6.  What  volume  of  ammonia  gas  will  be  liberated  by  the  action  of  any 
base  on  75  gm.  of  ammonium  sulphate?     (Standard  conditions.) 

7.  What  weight  of  ammonium  chloride  (95  per  cent  pure)  is  needed  for 
the  preparation  of  60  gm.  of  NH3?    Of  60  1.  at  22°  C.  and  767  mm.? 

8.  A  pupil  prepared  five  250  cc.  bottles  of  ammonia  gas  at  21°  C.  and 
755  mm.    What  weights  of  materials  interacted? 


CHAPTER  XI 

THE  ATMOSPHERE  —  ARGON  —  LIQUID  AIR 

120.  The  Atmosphere  is  the  vast  volume  of  gases  that 
envelops  the  earth  and  extends  many  miles  into  space. 
The  terms  atmosphere,  the  air,  and  air  are  often  used  inter- 
changeably. 

Aristotle  (384-322  B.  C.)  regarded  air  as  one  of  the  four  elementary 
principles  whose  combinations  made  up  all  substances  in  the  uni- 
verse. The  other  three  were  earth,  fire,  and  water.  He  taught 
that  air  possesses  two  fundamental  properties  —  heat  and  damp- 
ness. The  early  chemists  used  the  word  air  in  the  sense  in  which 
the  word  gas  is  now  employed.  Thus,  we  have  already  learned 
that  hydrogen  was  first  called  inflammable  air. 

121.  Atmospheric  Pressure.  --  This  enormous  mass  of 
gas  exerts  a  pressure  on  the  earth's  surface  called  atmos- 
pheric pressure,  which  is  about  fifteen  pounds  on  every 
square  inch.    The  amount  of  pressure  on  a  square  inch 
is  often  called  "an  atmosphere/'  and  it  is  sometimes  used 
as  a  unit  of  pressure.    Thus,   three  atmospheres  means 
a  pressure  of  forty-five  pounds  per  square  inch.    Atmos- 
pheric pressure  is  measured  by  an  instrument  called  a 
barometer,  and  the  pressure  at  a  given  time  is  found 
simply  by  reading  the  height  of  the  mercury  column  of 
the  barometer.      The  normal  or  standard  pressure  of  the 
atmosphere  is  equal  to  the  pressure  of  a  column  of  mer- 


THE   ATMOSPHERE  105 

cury  which  is  760  millimeters  (or  29.92  inches)  high. 
The  weight  of  a  liter  of  dry  air  at  o°  C.  and  760  mm.  is 
1.293  gm- 

122.  Ingredients   of   the   Atmosphere.  —  The   atmos- 
phere is  a  mixture  of  several  gases.     Oxygen,  nitrogen, 
and   argon   are   the   three   ingredients   that   are   always 
present  in  nearly  constant  proportions.     Variable  pro- 
portions of  water  vapor  and  carbon  dioxide  gas  are  always 
found,  and  also  small  quantities  of  compounds  related 
to  ammonia  and  nitric  acid.    Near  cities  the  atmosphere 
may  contain  considerable  dust,  sulphur  compounds,  and 
acids;    in   the   country  ozone  is  usually  present,  and  at 
the  ocean  some  salt  is  often  found. 

123.  Air  is  a  Mixture.  —  Chemical  compounds  have 
a  constant  composition,  i.e.  the  constituents  are  united 
in  a  proportion  which  is  always  the  same  in  the  case  of 
a  given  compound    (59).       Furthermore,  we  have  seen 
that  chemical   action  results  in  the  formation  of  one  or 
more  new  compounds,  and   that   this  action  is  usually 
accompanied  by  heat  changes.     The  following  facts  show 
that  air  is  a  mixture  of  free  gases :  - 

(1)  The  proportion  of  oxygen  and  of  nitrogen  is  not 
fixed,  but  varies  between  small  limits. 

(2)  When  nitrogen  and  oxygen  are  mixed  in  the  pro- 
portions that  form  air,  the  product  is  exactly  like  air,  but 
the  act  of  mixing  gives  no  evidence  of  chemical  action. 

(3)  When  air  is  dissolved  in  water,  a  larger  proportion 
of  oxygen  than  nitrogen  dissolves.     If  the  oxygen  and 
nitrogen  were  combined,  the  dissolved  air  would  contain 
the   same   proportions  of    oxygen    and    nitrogen   as   air 
itself. 

124.  Proportions  of  the  Constant  Ingredients  of  Air.  - 
For  many  years  it  was  believed  that  pure  air  consisted 


106  CHEMISTRY 

solely  of  oxygen  and  nitrogen.  But  in  1894  it  was  found 
that  nearly  1.2  per  cent  (by  volume)  of  the  gas  hitherto 
called  nitrogen  is  argon  (129).  The  normal  proportions 
(by  volume)  of  the  constant  ingredients  of  air  are  nitro- 
gen 78.122,  oxygen  20.941,  argon  .937. 

125.  Volumetric  Composition  of  Air.  —  Although  air  is  a  mixture, 
we  usually  speak  of  its  "  composition."     However,  the  proportions 
of  the  main  ingredients  of  air  are  so  nearly  constant,  chemists  have 
fallen  into  the  habit  of  applying  the  term  composition  to  air.     The 
proportion  of  oxygen  in  the  air  can  be  found  by  several  methods. 
In  one,  a  known  volume  of  air  is  shaken  in  a  closed  bottle  with  a 
mixture  of  pyrogallic  acid  and  potassium  hydroxide;   this  solution 
absorbs  the  oxygen  and  leaves  the  nitrogen  and  argon  unchanged. 
(See  Part  H,  Exp.  57.)      In  another,  a  graduated  glass  tube,  con- 
taining a  known  volume  of  air  is  inverted  in  a  jar  of  water,  and  a 
piece   of    phosphorus    attached  to  a  wire  is  introduced   into  the 
tube.      White  fumes  of  phosphorus  pentoxide  indicate  immediate 
action.      They  soon  dissolve  in  the  water,  which  rises  higher    in 
the  tube,  as  the  oxygen  combines  with  the  phosphorus.      In  a  few 
hours  the  phosphorus  is  removed,  and  the  volume  of  gas  is  read. 
The  difference  between  the  first  and  last  volumes  is  oxygen. 

126.  Oxygen  and  Nitrogen  in  the  Atmosphere.  --  The 
chemical  activity  of  the  atmosphere  is  due  to  the  free 
oxygen  it  contains.     Nitrogen  is  inactive,  and  if  the  at- 
mosphere contained  much  more  than  the  normal  amount, 
the  chemical  activity  of  the  oxygen  would  be  too  much 
retarded.     To  be  serviceable  to  man,  oxygen  must  be 
accompanied  by  the  proper  proportion  of  nitrogen. 

127.  Water  Vapor  in  the  Atmosphere.  —  Water  vapor 
is  always  present  in  the  atmosphere,  owing  to  constant 
evaporation  from  the  ocean  and  other  bodies  of  water. 
When  the  temperature  of  the  atmosphere  falls,  the  water 
vapor  condenses  and  is  deposited  in  the  form  of  dew, 
rain,  fog,  mist,  frost,  snow,  sleet,  or  hail.    The  clouds  are 


THE  ATMOSPHERE  107 

masses  of  water  vapor  which  have  been  condensed  by  the 
cold  upper  air.  A  given  volume  of  air  absorbs  a  definite 
volume  of  water  vapor  and  no  more.  Warm  air  holds 
more  vapor  than  cool  air.  Air  containing  its  maximum 
amount  of  water  vapor  is  said  to  be  saturated  at  that 
temperature,  or  to  contain  100  per  cent  of  water  vapor. 
The  saturation  point  is  also  called  the  dew  point.  On  a 
pleasant  day  in  a  temperate  climate  the  relative  humidity, 
i.e.  the  relative  amount  of  water  vapor  present,  may  vary 
from  30  to  90  per  cent. 

The  presence  of  water  vapor  in  the  air  is  shown  by  the  mois- 
ture which  collects  on  the  outside  of  a  vessel  containing  cold 
water,  such  as  a  pitcher  of  iced  water.  The  moisture  conies  from 
the  air  around  the  vessel.  For  a  similar  reason,  water  pipes  in  a 
cellar  and  the  cellar  walls  themselves  are  moist  in  summer.  The 
deliquescence  of  calcium  chloride,  common  salt,  and  other  substances 
likewise  reveals  the  presence  of  water  vapor  in  air  (51). 

128.  Carbon  Dioxide  in  the  Atmosphere.  —  Carbon 
dioxide  is  one  product  of  the  respiration  of  animals  and  of 
the  combustion  and  decay  of  organic  substances.  By  these 
processes  vast  quantities  of  carbon  dioxide  are  being  con- 
stantly introduced  into  the  atmosphere.  The  quantity  in 
the  atmosphere  is  variable,  though  not  between  such  wide 
limits  as  the  water  vapor.  The  proportion  in  ordinary 
air  is  3  to  4  parts  in  10,000  parts  of  air.  In  crowded  rooms 
it  is  often  as  high  as  33  parts  in  10,000.  The  proportion 
of  carbon  dioxide  in  the  atmosphere  as  a  whole  is  prac- 
tically constant,  largely  owing  to  the  fact  that  this  gas 
is  an  essential  food  of  plants  (186). 

The  presence  of  carbon  dioxide  in  air  is  detected  by  calcium 
hydroxide.  If  calcium  hydroxide  solution  is  exposed  to  air,  the 
carbon  dioxide  interacts  with  the  calcium  hydroxide,  forming  a 


io8  CHEMISTRY 

thin,  white  crust  of  insoluble  calcium  carbonate  on  the  surface  of 
the  liquid.  If  considerable  air  is  drawn  through  the  calcium  hy- 
droxide solution,  the  liquid  becomes  milky,  because  the  particles 
of  calcium  carbonate  are  suspended  in  the  liquid.  The  equation 
for  the  interaction  of  carbon  dioxide  and  calcium  hydroxide  is  — 

CO2  +        Ca(OH)2  CaC03         +     H2O 

Carbon  Dioxide          Calcium  Hydroxide         Calcium  Carbonate          Water 

Argon 

129.  Argon  in  the   Atmosphere.  —  Argon,   as   stated 
above,   is   an   essential   and   constant  ingredient   of   the 
atmosphere,  the  proportion  being  .937  per  cent  by  vol- 
ume.    Argon  was  detected  and  first  studied  in  1894  by 
Rayleigh  and  Ramsay.     Rayleigh  found   that  nitrogen 
extracted  from  air  had  a  greater  weight  than .  an  equal 
volume  of  nitrogen  obtained  from  compounds  of  nitro- 
gen.   Consequently,  they  believed  that  the  nitrogen  from 
air  contained  another  gas  hitherto  overlooked.     Experi- 
ments showed  that  after  the  oxygen  and  nitrogen  were 
removed  from  purified  air,  there  still  remained  a  small 
quantity  of  a  new  gas.    They  named  it  argon  and  gave  it 
the  symbol  A. 

Argon  can  be  obtained  by  passing  pure  air  over  heated  copper 
to  remove  the  oxygen,  and  then  the  remaining  gas  over  heated 
magnesium  or  calcium  to  remove  the  nitrogen.  Another  method 
consists  in  passing  electric  sparks  through  a  mixture  of  air  and 
oxygen,  and  absorbing  the  oxides  of  nitrogen  in  potassium  hydrox- 
ide solution.  The  latter  method  is  a  repetition  of  the  one  used 
by  Cavendish  in  1785  when  he  determined  the  composition  of  air. 
He  observed  and  recorded  the  fact  that  a  small  bubble  of  gas 
always  remained;  it  was  doubtless  argon,  and  to  Cavendish  belongs 
the  honor  of  first  observing  this  element. 

130.  Properties  of  Argon.  —  Argon  is  a  colorless,  odor- 
less gas  which  is  a  little  heavier  than  oxygen.     It  dis- 


LIQUID   AIR  109 

solves  in  water  to  the  extent  of  about  4  volumes  in  100. 
It  has  been  liquefied  and  solidified;  the  boiling  point  of 
liquid  argon  is  —  i86°C.  and  the  melting  point  of  the 
solid  (which  is  colorless)  is  — 189.5°  C.  A  conspicuous 
property  of  argon  is  its  lack  of  chemical  activity.  No 
compounds  of  this  element  have  as  yet  been  prepared  or 
discovered.  The  name  argon  is  happily  chosen,  being 
derived  from  Greek  words  signifying  inert. 

131.  Rare  Gases  in  the  Atmosphere.  —  Helium  (He),  neon  (Ne), 
krypton  (Kr),  and  xenon  (Xe)  are  inert  gases  discovered  by  Ram- 
say subsequently  to  argon.      With  the  exception  of  neon  they  con- 
stitute an  exceedingly  minute  proportion  of  the  atmosphere.      Like 
argon   they  do   not  form  compounds.      Ramsay  estimates  that  in 
1,000,000  parts  of  the  atmosphere  there  are  i  to   2  parts  of  hel- 
ium, 10  to  20  of  neon,  .05  of  krypton,  and  .006  of  xenon.    Helium 
was  detected  in  the  atmosphere  of  the  sun  by  Lockyer  in  1868. 
It  was  found  by  Ramsay,  soon  after  he  discovered  argon,  in  the 
gases  expelled  from  certain  rare  minerals  and  in  the  gas  and  water 
of  some  mineral  springs.      Recently  it  has  become  conspicuous  as 
one  of  the  disintegration  products  of  radium  (529). 

Liquid  Air 

132.  Liquid   air  is   a  mixture  of   the  liquefied  gases 
that  constituted  the  air  used.     It  is  a  milky  liquid,  owing 
to  the  presence  of  solid  carbon  dioxide  and  ice.    If  these 
solids  are  removed  by  filtering,  the  filtrate  has  a  pale 
blue  tint.     It  boils  at  about  —190°  C.  under  atmospheric 
pressur-e.     If  an  ordinary  vessel  is  filled  with  liquid  air, 
the  latter  boils  vigorously,  the  surrounding  air  becomes 
intensely  cold,  frost  gathers  on  the  vessel,  and  in  a  short 
time  the  liquid  air  will  have  entirely  disappeared  into  the 
air  of  the  room.     If,  however,  liquid  air  is  placed  in  a 
Dewar  flask,  evaporation  takes  place  so  slowly  that  some 
liquid  air  will  remain  in  the  flask  several  days. 


no  CHEMISTRY 

A  Dewar  flask  (Fig.  31)  consists  of  two  flasks,  one  within  the 
other,  sealed  together  air-tight  at  the  top;  the  space 
between  the  flasks  is  a  vacuum.  The  surfaces  of 
the  flasks  are  coated  with  silver,  which  reflects  heat 
and  helps  retard  the  evaporation  of  the  liquid  air. 
Liquid  air  is  stored  and  transported  in  Dewar  flasks. 

133.   Changes  produced  by  Liquid  Air.  - 

Liquid  air,  owing  to  its  extremely  low  tem- 
perature, produces  remarkable  physical 
changes.  A  tin  or  iron  vessel  which  has 
been  cooled  by  liquid  air  is  so  brittle  that  it 
Dewar  flask  may  °^en  ^e  crushed  with  the  fingers.  Mer- 
cury freezes  so  hard  in  liquid  air,  that  it  may 
be  used  as  a  hammer  to  drive  a  nail.  When  liquid  air  is 
put  in  a  teakettle  standing  on  a  block  of  ice,  the  liquid  air 
boils  vigorously.  If  the  kettle  of  liquid  air  is  placed  over 
alighted  Bunsen  burner,  frost  and  ice  collect  on  the  bot- 
tom of  the  kettle,  because  the  intense  cold  produced  by 
the  evaporation  of  the  liquid  air  in  the  kettle  solidifies 
the  water  vapor  and  carbon  dioxide,  which  are  the  two 
main  products  of  burning  illuminating  gas.  If  water  is 
now  poured  into  the  kettle,  the  liquid  air  boils  over  and 
the  water  is  instantly  frozen;  the  water  is  so  much  hotter 
than  the  liquid  air  that  the  latter  boils  more  violently, 
and  since  its  rapid  evaporation  causes  the  absorption  of 
heat,  the  water  gives  up  its  heat  and  becomes  ice.  Ordi- 
nary liquid  air  is  from  one  half  to  one  fifth  liquid  oxygen, 
and  will  support  combustion.  A  red-hot  rod  of  steel  or 
of  carbon  burns  brilliantly  in  this  cold  liquid. 

134.  Numerous  applications  of  liquid  air  have  been  proposed, 
and  some  have  passed  the  experimental  stage.  It  is  used  for  re- 
moving diseased  flesh  from  a  wound,  and  as  a  commercial  source 
of  oxygen  and  nitrogen.  The  last  use  is  based  primarily  on  the 


LIQUID   AIR  in 

fact  that  when  liquid  air  evaporates,  the  nitrogen  passes  off  first, 
and  in  a  short  time  relatively  pure  oxygen  remains. 

135.  Liquid  air  is  manufactured  in  large  quantities  at  a  com- 
paratively low  cost.  Compressed  air  cooled  by  water  is  forced 
through  a  pipe  to  a  valve.  As  it  escapes  through  the  valve,  it 
expands  and  its  temperature  falls,  because  expansion  is  a  cooling 
process.  After  expansion  the  cold  air  is  led  back  over  the  outer 
surface  of  the  same  pipe  by  which  it  came,  whereupon  it  rapidly 
regains  its  former  temperature.  But  in  doing  so  it  cools  the  pipe 
itself  and  the  air  within  it.  This  latter  air  in  turn  expands  and 
falls  in  temperature,  but  as  it  was  colder  than  the  first  portion 
before  expansion,  so  it  is  colder  after  expansion.  Since  the  pres- 
sure within  the  pipe  is  maintained  by  a  continuous  supply  of  com- 
pressed air,  the  pipe  becomes  continually  colder  until  finally  the 
expanding  air  at  the  valve  liquefies  in  part  and  is  collected  in  a 
suitable  receptacle. 

EXERCISES 

1.  What  are  the  two  chief  ingredients  of  the  atmosphere?    The  perma- 
nent ingredients?     The  variable  ingredients?     The  ingredients  found  in 
traces?    What  special  substances  are  sometimes  found  in  the  air  of  cities? 

2.  Compare  the  functions  of  oxygen  and  nitrogen  in  the  atmosphere. 

3.  State  the  volumetric  composition  of  air.    Has  air  a  chemical  formula? 
If  so,  what  is  it?     If  not,  why? 

4.  Describe  the  action  of  air  upon  (a)  calcium  hydroxide  and  (6)  cal- 
cium chloride. 

6.   Give  several  proofs  that  air  is  a  mixture. 

6.  What  is  argon?      Give  a  brief  account  of  (a)  its  discovery,  (6)  its 
properties,  (c)  its  method  of  preparation.    What  proportion  of  air  is  argon? 
What  is  the  significance  of  the  name  argon? 

7.  What  is  liquid  air?    What  are  its  chief  properties?    State  briefly  its 
method  of  manufacture.    Describe  a  Dewar  flask. 

PROBLEMS 

1.  What  is  the  weight  of  air  in  a  room,  6  X  8  X  5  m.,  if  a  liter  of  air  weighs 
1.293  gm.? 

2.  How  many  kilograms  of  pure  air  are  needed  to  yield  (a)  100  kg.  and 
(6)  100  1.  of  oxygen?    (Standard  conditions.) 

3.  Express  in  inches  the  following  barometer  readings:    (a)  760  mm., 
(6)  745  mm.,  (c)  70  cm.,  (d)  0.769  m.,  (e)  7.49  dm.,  (/)  780  mm.,  (g)  5  mm. 


ii2  CHEMISTRY 

4.  What  is  the  weight  at  o°  C.  and  760  mm.  of  (a)  1000  cc.  of  dry  air? 
Of  (b)  95  1.,  (c)  95  cc.,  (d)  95  cu.  m.? 

5.  One  liter  of  air  under  standard  conditions  weighs  1.293  8m-    What  is 
the  weight  of  2.5  liters  when  the  barometer  stands  at  755  mm.? 

6.  What  volume  would  10  1.  of  air  at  25°  C.  occupy  at  o°  C.?     (Pres- 
sure unchanged.) 

7.  How  many  cc.  will  i  gm.  of  air  occupy?     (Standard  conditions.) 

8.  A  flask  weighed  130  grams  when  full  of  air,  and  129.84  grams  when 
some  of  the  air  was  sucked  out.    When  opened  under  water  125  cc.  of  water 
entered.    Find  the  weight  of  a  liter  of  air. 


CHAPTER  XII 

GAY-LUSSAC'S     LAW     OF     GAS     VOLUMES  —  AVOGADRO'S 

HYPOTHESIS  —  MOLECULAR  WEIGHTS  AND  ATOMIC 

WEIGHTS  — MOLECULAR  FORMULAS  AND  EQUATIONS 

Atomic  weights  and  molecular  weights  have  been  used 
freely  in  the  foregoing  pages.  In  the  present  chapter, 
after  discussing  Gay-Lussac's  law  and  Avogadro's  hy- 
pothesis, we  shall  consider  the  methods  by  which  these 
weights  are  determined. 

136.  Gay-Lussac's  Law  of  Gas  Volumes. — We  have 
already  seen  that  gases  combine  by  volume  in  simple 
ratios.  These  results  may  be  summarized  in  a  - 

TABLE  OF  THE  COMBINATION  OF   GASES  BY  VOLUME 


Volumes  of  Combining  Gases 

Volumes  of  Gaseous  Product 

2  volumes  of  hydrogen 
i  volume  of  oxygen 

2  volumes  of  water  vapor 

i  volume  of  chlorine 
i  volume  of  hydrogen 

2  volumes  of  hydrogen  chloride 

3  volumes  of  hydrogen 
i  volume  of  nitrogen 

2  volumes  of  ammonia  gas 

2  volumes  of  nitrogen 
i  volume  of  oxygen 

2  volumes  of  nitrous  oxide  gas 

i  volume  of  nitrogen 
i  volume  of  oxygen 

2  volumes  of  nitric  oxide  gas 

H4  CHEMISTRY 

It  is  clear  from  the  above  table  that  in  the  case  of  these 
gases  small  whole  numbers  express  the  relation  existing 
between  the  volumes  of  the  combining  gases  and  the 
volume  of  the  gaseous  product.  The  simple  ratio  that 
exists  between  the  gas  volumes  (tabulated  above),  whether 
components  or  products,  is  true  of  all  gases.  This  general 
fact  was  summarized  in  1808  by  the  French  chemist  Gay- 
Lussac  in  the  form  of  a  law,  thus :  - 

Gases  combine  in  volumes  which  bear  a  simple  numeri- 
cal ratio  to  each  other  and  to  the  volume  of  their  gaseous 
product. 

By  a  " simple  numerical  ratio"  is  meant  one  involving 
only  small  whole  numbers. 

137.  Avogadro's    Hypothesis.  —  In    1811    the   Italian 
physicist  Avogadro  proposed  an  explanation  of  the  simple 
numerical  relation  of  gas  volumes.     It  is  usually  called 
Avogadro's  hypothesis  and  may  be  stated  thus :  - 

Equal  volumes  of  all  gases  at  the  same  temperature  and 
pressure  contain  approximately  the  same  number  of  molecules. 

This  statement  means,  for  example,  that  a  liter  of 
oxygen,  a  liter  of  hydrogen  chloride,  of  nitric  oxide,  or  of 
any  other  gas,  at  the  same  temperature  and  pressure, 
contains  very  nearly  the  same  number  of  molecules. 

138.  Relative  Weights  of  Molecules.  —  By  means  of 
Avogadro's  hypothesis  we  can  find  the  relative  weights  of 
molecules.     Let  us  take  the  case  of  oxygen  and  carbon 
dioxide.    A  liter  of  carbon  dioxide  weighs  1.977  gm.  and 
a  liter  of  oxygen  1.429  gm.  at  the  same  temperature  and 
pressure,  i.e.  the  weight  of  a  liter  of  carbon  dioxide  is 
approximately  1.38  times  that  of  a  liter  of  oxygen.    Hence 
according  to  Avogadro's  hypothesis  the  weight  of  the 
carbon  dioxide  molecules  is  about  1.38  times  the  weight  of 
the  oxygen  molecules. 


MOLECULAR   AND   ATOMIC   WEIGHTS       115 

139.   Determination  of  Approximate  Molecular  Weights. 

—  Approximate  molecular  weights  of  gases  or  of  vola- 
tilized substances  are  determined  by  two  steps,  (i)  find- 
ing by  experiment  the  vapor  density  referred  to  oxygen, 
and  (2)  multiplying  this  value  by  32.  The  expression 
vapor  density  referred  to  oxygen,  as  used  here,  means  the 
number  found  by  dividing  the  weight  of  a  given  volume  of 
a  gas  or  vapor  by  the  weight  of  an  equal  volume  of  oxy- 
gen (measured  at  the  same  temperature  and  pressure). 
Thus,  in  the  example  given  in  the  preceding  paragraph  the 
number  1.38  is  the  vapor  density  of  carbon  dioxide.  And 
since  this  number,  as  we  have  seen,  expresses  the  relation 
of  the  weights  of  the  carbon  dioxide  and  oxygen  molecules, 
it  is  evident  that  we  could  find  the  weight  of  a  molecule 
of  carbon  dioxide  if  we  knew  the  weight  of  a  molecule  of 
oxygen.  Now  the  weight  of  a  molecule  of  oxygen  is  32, 
and  it  is  for  this  reason  that  the  vapor  density  is  multi- 
plied by  32  in  finding  molecular  weights. 

The  weight  of  a  molecule  of  oxygen  is  32,  because  a 
molecule  of  oxygen  contains  two  atoms  each  having  the 
weight  1 6.  The  conclusion  that  a  molecule  of  oxygen 
contains  two  atoms  is  based  mainly  on  the  following  argu- 
ment, which  involves  Gay-Lussac's  law  and  Avogadro's 
hypothesis :  - 

When  oxygen  and  nitrogen  combine  to  form  nitric 
oxide,  one  volume  of  oxygen  combines  with  one  volume 
of  nitrogen  to  form  two  volumes  of  nitric  oxide.  Sup- 
pose the  volume  of  oxygen  contains  100  molecules. 
Then,  according  to  Avogadro's  hypothesis,  the  equal 
volume  of  nitrogen  contains  100  molecules,  while  the  two 
volumes  of  the  product  contain  200  molecules  of  nitric 
oxide.  That  is:  — 


n6  CHEMISTRY 

100  molecules  of  Oxygen  +  100  molecules  of  Nitrogen 

=  200  molecules  of  Nitric  Oxide. 

Now,  since  every  molecule  of  nitric  oxide  contains  at  least 
one  atom  each  of  oxygen  and  nitrogen,  the  200  mole- 
cules must  contain  at  least  200  atoms  of  oxygen.  But  the 
200  atoms  of  oxygen  were  provided  by  the  100  molecules 
of  oxygen.  Therefore,  each  molecule  of  oxygen  must 
contain  at  least  two  atoms.  There  is  conclusive  evi- 
dence (based  on  certain  physical  properties  of  gases 
when  heated)  that  the  oxygen  molecule  contains  only 
two  atoms. 

The  method  of  determining  the  approximate  molecular 
weight  of  a  substance  is  now  clear.  It  is  only  necessary 
to  find  the  vapor  density  on  the  oxygen  basis  and  mul- 
tiply this  value  by  32.  That  is:  — 

Molecular  Weight  =  Vapor  DensityVef  erred  to  Oxygen  X32. 

Thus,  since  the  vapor  density  of  carbon  dioxide  is  1.38, 
its  approximate  molecular  weight  is  1.38  X  32,  or  44.16. 

Some  substances  cannot  be  vaporized  without  decomposition. 
The  molecular  weights  of  such  substances  cannot,  of  course,  be 
found  by  the  vapor  density  method.  If  a  substance  dissolves 
without  decomposition,  the  molecular  weight  of  the  dissolved  sub- 
stance can  be  determined  by  an  appropriate  method  (162).  No 
experimental  method  is  known  for  determining  the  molecular 
weight  of  a  substance  in  the  solid  state  (i.e.  not  dissolved  or 
vaporized);  it  is  customary  to  assume  that  the  molecular  weight 
of  such  substances  is  the  sum  of  the  atomic  weights  in  the  simplest 
formula. 

140.   Determination  of  Approximate  Atomic  Weights.  - 
The  approximate  atomic  weight  of  an  element  is  deter- 
mined  from   the   molecular   weights   of  its   compounds. 
We  have  already  seen  that  molecular  weights  can  be  found 
by   multiplying   the   vapor   density   by   32.      Molecular 


MOLECULAR  AND   ATOMIC   WEIGHTS       117 

weights  thus  determined  are  approximate,  because  Avo- 
gadro's  hypothesis  is  approximate;  that  is,  they  are  not 
exactly  (though  often  very  nearly)  equal  to  the  sum  of 
the  exact  weights  of  the  atoms  in  one  molecule.  For 
example,  the  approximate  molecular  weight  of  carbon 
dioxide  is  44.16  (139),  whereas  the  exact  weight  is  44.00. 
Hence  atomic  weights  derived  from  molecular  weights -are 
approximate.  Subsequently  the  method  of  finding  exact 
atomic  weights  will  be  discussed  (141). 

After  the  molecular  weights  of  compounds  have  been 
determined,  the  next  problem  is  to  find  what  parts  of  the 
molecular  weights  should  be  chosen  as  the  atomic  weights 
of  the  elements  that  constitute  the  compounds.  The 
steps  in  the  procedure  are:  First,  determine  the  molec- 
ular weights  of  several  compounds  of  these  elements; 
second,  find  by  analysis  the  per  cent  of  each  element  in 
the  compounds;  third,  find  the  weight  of  each  element 
in  this  molecular  weight  by  multiplying  the  molecular 
weight  of  each  compound  by  the  per  cent  of  the  elements 
in  the  compound.  The  minimum  value  obtained  in  the 
case  of  each  element  will  be  the  approximate  atomic 
weight.  The  numerical  results  obtained  from  a  study 
of  the  elements  oxygen,  hydrogen,  chlorine,  nitrogen,  and 
carbon  are  shown  in  the  table  on  page  118,  in  which,  for 
the  sake  of  simplicity,  whole  numbers  are  used  (except 
in  the  case  of  chlorine) . 

In  this  table,  columns  one  and  two  contain  the  names  of 
the  compounds  and  their  approximate  molecular  weights. 
The  other  columns  contain  the  parts  of  the  molecular 
weights  that  belong  to  the  atoms  of  the  elements  in 
a  molecule  of  the  compound.  These  values  are  obtained 
by  the  third  step  in  the  procedure.  For  example,  the  pro- 
cedure in  the  case  of  water  is  as  follows:  (i)  by  experi- 


n8 


CHEMISTRY 


DETERMINATION  OF  APPROXIMATE  ATOMIC  WEIGHTS 


Compound 

Molecular 
Weight 

Weight 
of 
Oxygen 

-4->             <D 

•S^* 
Z°-v 

*  6 

Weight 
of 
Chlorine 

14 
*  g 

Weight 
of 
Carbon 

Water  

18 
34 
36.5 
17 
63 
44 
30 
46 
28 

44 
16 
28 
26 

74 
46 

iiQ-5 
154 
61.5 

16 

32 

48 
16 
16 

32 

16 
32 

16 
16 

2 
2 

I 

3 

i 

4 
4 

2 
IO 

6 

i 

35-5 

106.5 
142 
35-5 

14 
14 
28 

14 
14 

14 

12 
12 
12 
24 
24 
48 
24 
12 
12 
12 

Hydrogen  Dioxide 

Hydrogen  Chloride  
Ammonia  

Nitric  Acid 

Nitrous  Oxide   . 

Nitric  Oxide  

Nitrogen  Dioxide  
Carbon  Monoxide  

Carbon  Dioxide  

Methane  

Ethylene 

Acetylene  

Ether 

Ethyl  Alcohol  . 

Chloroform 

Carbon  Tetrachloride  .  .  . 
Cyanogen  Chloride  

Minimum  weight  of  each  element.  . 

16 

i 

35-5 

14 

12 

ment  the  approximate  molecular  weight  of  water  vapor 
is  found  to  be  18;  (2)  by  analysis  the  compound  is  shown 
to  contain  88.82  per  cent  oxygen  and  11.18  per  cent 
hydrogen;  and  (3)  the  products  of  18  and  these  per- 
centages are  16  and  2  respectively.  An  examination  of 
the  weights  of  the  elements  show  (i)  that  the  minimum 
weight  in  each  column  is  O  =  16,  H  =  i,  Cl  =  35.5, 
N  =  14,  and  C  =  12,  and  (2)  that  the  other  weights  are 
simple  multiples.  These  facts  are  significant.  The  small- 
est weights  must  be  the  weights  of  single  atoms,  for  it  is 
highly  probable  that  one  or  more  compounds  in  a  repre- 


MOLECULAR  AND   ATOMIC   WEIGHTS       119 

sentative  group  will  contain  only  one  atom  of  a  given 
element,  and  the  part  of  the  molecular  weight  appor- 
tioned to  the  element  in  question  will  of  course  be  its 
atomic  weight.  In  the  compounds  that  contain  this  ele- 
ment in  a  multiple  proportion,  it  is  likewise  obvious  that 
the  molecule  must  contain  several  atoms  of  the  element. 
Thus,  in  hydrogen  dioxide  the  weight  of  oxygen  is  twice 
that  in  water,  and  we  conclude  that  a  molecule  of  hydro- 
gen dioxide  contains  two  atoms  of  oxygen  —  a  conclusion 
in  harmony  with  other  observations. 

Formerly  the  selection  of  the  atomic  weights  of  solid  elements, 
especially  metals,  was  checked  by  an  approximate  generalization 
commonly  called  the  law  of  Dulong  and  Petit,  which  can  be  stated 
thus:  — 

The  atomic  weight  is  the  quotient  of  6.25  (approximately)  divided 
by  the  specific  heat. 

More  reliable  methods  are  now  used  in  determining  atomic 
weights,  although  this  so-called  law  is  helpful  in  deciding  between 
a  number  and  its  multiple. 

141.   Accurate  Determination  of  Atomic  Weights.  — 

The  methods  of  determining  atomic  weights  discussed 
in  the  preceding  sections  yield  approximate  values. 
When  the  approximate  value  is  found,  the  accurate  value 
is  determined  by  painstaking  analysis  of  very  carefully 
purified  substances.  The  general  method  can  be  illustrated 
by  an  actual  case  kindly  furnished  by  the  American 
chemist  Richards,  who  has  made  marvellously  accurate  de- 
terminations of  atomic  weights.  He  found  that  28.26299 
gm.  of  silver  chloride  were  formed  from  21.27143  gm.  of 
silver.  He  accepted  AgCl  as  the  formula  of  silver  chloride 
and  107.880  as  the  atomic  weight  of  silver,  and  calculated 
the  atomic  weight  of  chlorine  thus:  - 


120  CHEMISTRY 

28.26299  —  21.27143  =  6.99156 

Wt.  of  silver  :  Wt.  of  chlorine  ::  At.  wt.  of  silver  :  At.  wt.  of  chlorine 
21.27143  :        6.99156        ::  107.880        :  x 

x  =  35458 

The  international  atomic  weight  (35.46)  is  based  on  this 
and  other  determinations  made  by  the  same  chemist. 

The  exact  determination  of  atomic  weights  is  a  difficult 
task.  And  although  the  utmost  care  is  used  in  purify- 
ing the  chemicals  and  performing  the  analysis,  the  results 
of  different  experimenters  do  not  always  exactly  agree. 
Therefore  an  international  committee  was  chosen  several 
years  ago  to  select  the  most  accurate  atomic  weights  of 
the  elements.  These  weights  are  embodied  in  a  table 
published  annually  and  called  the  International  Table 
of  Atomic  Weights.  The  table  is  given  on  the  inside  of 
the  back  cover  of  this  book.  In  this  table  the  accepted 
atomic  weights  are  placed  in  one  column  and  the  ap- 
proximate values  in  another.  The  approximate  atomic 
weights  are  sufficiently  accurate  for  general  reference  and 
in  making  chemical  calculations;  they  may  be  used  in 
solving  the  problems  in  this  book. 

142.  Formulas  of  Molecules  of  Compounds.  —  A 
formula  expresses  composition,  that  is,  it  represents  by 
means  of  symbols  the  kind  and  number  of  atoms  in  a 
molecule.  This  means,  of  course,  that  the  sum  of  the 
weights  represented  by  the  atoms  in  a  molecule  must 
equal  approximately  the  molecular  weight  of  the  substance 
found  by  experiment.  Let  us  take  two  illustrations, 
(i)  Analysis  shows  that  the  ratio  of  hydrogen  to  chlorine 
in  hydrogen  chloride  is  i  to  35.5,  and  a  supplementary 
experiment  shows  that  the  molecular  weight  of  hydrogen 
chloride  is  approximately  36.7.  Therefore,  the  formula  of 
a  molecule  of  hydrogen  chloride  is  HC1,  first,  because  this 


MOLECULAR  AND   ATOMIC   WEIGHTS       121 

expression  represents  the  ratio  found  by  analysis  (i  :  35.5), 
and  second,  because  the  molecular  weight  found  by  experi- 
ment (36.7)  is  approximately  the  sum  of  the  weights  of 
the  atoms  (i  +  35.5  =  36.5)  in  a  molecule.  (2)  Experi- 
ment shows  that  the  ratio  of  hydrogen  to  oxygen  in  hydro- 
gen dioxide  is  i  to  16,  and  that  the  molecular  weight  is 
approximately  34.  Now  HO  would  correctly  express  the 
composition  of  hydrogen  dioxide  as  found  simply  by 
analysis,  because  H  and  O  represent  the  ratio  i  to  16. 
But  the  molecular  weight  represented  by  HO  is  17  (i.e. 
i  +  1 6),  which  is  only  half  of  the  approximate  molec- 
ular weight  found  by  experiment.  Therefore  H2O2,  not 
HO,  must  be  the  correct  molecular  formula  of  hydrogen 
dioxide,  because  H202  expresses  not  only  the  ratio  (i  :  16) 
but  also  the  molecular  weight  (34). 

In  a  previous  section  (68  (i))  it  was  shown  that  the 
simplest  formula  of  a  compound  can  be  readily  calculated 
from  the  percentage  composition,  viz.  by  dividing  the 
per  cent  of  each,  element  in  the  compound  by  the  'atomic 
weight  (and,  if  necessary,  reducing  the  quotients  to  the 
smallest  whole  numbers).  That  is,  we  can  transform  the 
percentages  found  by  analysis  into  numbers  which  express 
the  atomic  relations  by  distributing  the  proportional  parts 
among  the  elements  according  to  that  method  of  express- 
ing composition  adopted  in  chemistry.  Formulas  thus 
calculated,  however,  are  not  necessarily  molecular  formu- 
las, (i)  If  the  molecular  weight  of  a  compound  cannot  be 
found  by  experiment,  then  the  simplest  formula  is  ac- 
cepted as  the  molecular  formula.  For  example,  according 
to  analysis,  a  compound  contains  40  per  cent  of  calcium, 
12  of  carbon,  and  48  of  oxygen.  Dividing  each  per  cent 
by  the  proper  atomic  weight,  we  have:  —  40  ^40  =  i, 
12  -J-  12  =  i,  48  -T-  16  =  3.  That  is,  one  molecule  of 


122  CHEMISTRY 

this  compound  contains  (at  least)  one  atom  each  of  cal- 
cium and  carbon,  and  three  of  oxygen,  and  the  simplest 
formula  is  CaCOs.  This  is  also  accepted  as  its  molecular 
formula,  because  the  molecular  weight  cannot  be  found 
by  any  method  known  at  present.  (2)  If,  however,  the 
molecular  weight  can  be  found  (by  means  of  the  vapor 
density),  the  molecular  formula  can  be  readily  calcu- 
lated. For  example,  a  compound  was  found  by  analysis 
to  contain  92.3  per  cent  of  carbon  and  7.7  of  hydro- 
gen, and  to  have  a  vapor  density  of  2.4375.  Proceed- 
ing as  above,  we  have  92.3  -f-  12  =  7.7,  and  7.7  -f- 1  = 
7.7.  Since  y.yiy.yasiii,  the  compound  contains  at 
least  one  atom  each  of  carbon  and  hydrogen,  and  would 
have  the  formula  CH,  if  nothing  were  known  about  its 
molecular  weight.  The  vapor  density  2.4375  requires 
the  molecular  weight  78  (i.e.  2.4375  X  32),  which  is  six 
times  the  weight  (13)  corresponding  to  the  formula  CH. 
Hence  the  molecular  formula  of  this  compound  is  not  CH, 
but  C6H6. 

To  recapitulate  regarding  molecular  formulas  of  com- 
pounds: The  simplest  formula  of  a  compound  is  found  by 
dividing  the  per  cent  of  each  element  by  its  atomic  weight 
and  reducing  these  quotients  to'  the  smallest  whole  num- 
bers (if  necessary);  if  the  molecular  weight  cannot  be 
found  by  experiment,  this  formula  is  accepted  as  the 
molecular  formula.  The  molecular  formula  of  a  com- 
pound is  found  by  three  steps:  (a)  Find  the  simplest 
formula,  (b)  divide  the  molecular  weight  by  the  sum  of 
the  weights  of  the  atoms  in  the  simplest  formula,  (c)  mul- 
tiply the  integral  numbers  of  the  simplest  formula  by  the 
quotient  obtained  in  (b). 

143.  Molecular  Weights  and  Molecular  Formulas  of 
Elements.  —  Several  elements  are  gases  at  ordinary  tem- 


'MOLECULAR  AND   ATOMIC   WEIGHTS       123 

peratures,  and  others  can  be  changed  into  vapors  by  heat- 
ing. The  molecular  weights  of  elements,  as  well  as  of 
compounds,  can  be  determined  by  finding  their  vapor 
densities  and  multiplying  these  values  by  32.  Such  de- 
terminations show:  (i)  The  gaseous  elements  already 
studied,  as  well  as  some  others,  have  molecular  weights 
which  are  twice  the  atomic  weight;  that  is,  the  molecule 
consists  of  two  atoms,  and  their  molecular  formulas  are, 
for  example,  O2,  H2,  C12,  N2  (compare  139).  (2)  The 
molecular  weights  of  most  metallic  elements  and  certain 
gaseous  elements  are  identical  with  their  atomic  weights; 
that  is,  the  molecule  consists  of  one  atom,  and  the  molec- 
ular formula  is  the  same  as  the  atomic  symbol,  e.g. 
Na,  K,  Zn,  Hg,  Cd  (cadmium),  A  (argon),  He  (helium), 
and  Ne  (neon) .  (3)  The  molecular  weights  of  certain  ele- 
ments vary  with  the  temperature,  decreasing  with  rise 
of  temperature;  e.g.  at  lower  temperatures,  molecules 
of  iodine,  sulphur,  and  phosphorus  are  represented  by 
I2,  Ss,  and  P4,  and  at  higher  temperatures  by  I,  S2,  and  P2 
(see  Physical  Properties  of  these  elements). 

In  representing  free  or  uncombined  elements,  molec- 
ular formulas  are  the  correct  chemical  expressions.  In 
the  succeeding  chapters  gases  will  be  represented  by  the 
proper  molecular  formulas,  unless,  of  course,  the  atomic 
symbols  answer  the  purpose;  the  elements  that  are  solid 
(or  liquid,  like  mercury)  at  ordinary  temperatures  will  be 
represented  by  atomic  expressions  (e.g.  Zn,  28,  4P)  except 
when  conditions  demand  molecular  formulas. 

144.  Molecular  Equations.  —  Hitherto  the  equation 
for  the  reaction  between  hydrogen  and  oxygen  has  been 
written:  - 

2H  +  O  =  H2O 


124  CHEMISTRY 

Since  the  reaction  is  between  gases  and  the  molecular 
formula  of  hydrogen  is  H2  and  of  oxygen  is  O2,  the  equation 
can  be  written:  - 

2H2  +  02  =  2H2O 

This  equation  is  read  thus:  Two  molecules  of  hydrogen 
unite  with  one  molecule  of  oxygen  to  form  two  molecules 
of  water  vapor.  Since  this  equation  correctly  represents 
the  interacting  substances  as  molecules,  the  equation 
is  called  a  molecular  equation.  It  should  be  noted  that 
the  proportions  by  weight  are  the  same  as  in  the  simpler 
or  atomic  form  of  the  equation.  Molecular  equations 
are  sometimes  called  volume  equations  or  gas  equations, 
because  they  show  the  volumes  of  gases  involved  in  the 
reaction.  Thus,  the  equation 

H2  +  C12  =  2HC1 
may  be  written  — 

H2      +     C12      =     2HC1 

i  volume         i  volume        2  volumes 

because  equal  numbers  of  molecules  represent  equal  vol- 
umes (137).  This  equation  is  read:  One  volume  of 
hydrogen  and  one  volume  of  chlorine  form  two  volumes 
of  hydrogen  chloride.  It  should  be  remembered,  then, 
that  in  molecular  or  volume  equations  a  single  molecule 
represents  one  volume  of  a  gas  (or  vapor). 

EXERCISES 

1.  State    and    illustrate    (a)    Gay-Lussac's   law  and    (6)   Avogadro's 
hypothesis. 

2.  What  is  the  relation  of  the  molecular  weight  of  a  gas  to  (a)  the 
molecular  and  (b)  the  atomic  weight  of  oxygen? 

3.  (a)  State  the  argument  proving  that  a  molecule  of  oxygen  con- 
sists of  at  least  two  atoms,     (b)  Apply  the  same  argument  to  nitrogen. 

4.  Hydrogen  and  nitrogen  combine  in  the  ratio  of  3  to  i  to  form  2  vol- 


MOLECULAR  AND   ATOMIC   WEIGHTS       125 

umes  of  ammonia.     Show  from  this  relation  that  a  molecule  of  nitrogen 
contains  at  least  two  atoms. 

6.  What  is  the  relation  between  molecular  weight  and  vapor  density? 
Illustrate  your  answer. 

6.  Why  is  the  formula  of  water  vapor  H2O  and  not  HO  or  H2O2? 

7.  How  are  molecular  weights  determined? 

8.  How  are  approximate  atomic  weights  found  from  molecular  weights? 

9.  What  is  a  molecular  formula?     What  is  the  molecular  formula  of 
oxygen,  nitrogen,  chlorine,  hydrogen,  zinc,  iodine,  sulphur,  mercury,  phos- 
phorus, sodium,  potassium,  argon,  helium?     How  is  a  molecular  formula 
determined?     Illustrate  your  answer. 

10.  Write  the  chemical  expression  for:   (a)  one  atom,  three  atoms,  one 
molecule,  and  three  molecules  of  hydrogen,  of  .zinc,  and  of  phosphorus, 
and  (b)  one  molecule  of  water,  three  molecules  of  water. 

11.  Explain  fully:    O,  O2,  O3,  3O2,  2O3,  KC1O3,   4P,    2P,  P4,  P2,  2?4, 
4P2,  P205,  Hg,  HgO,  Zn,  2Zn,  Na,  2Na,  Na2O,  Na3AlO3,  C,  CO,  CO2,  CiaHaOn. 

12.  Express  the  following  by  an  equation:   One  volume  of  phosphorus 
vapor  and  six  volumes  of  chlorine  form  four  volumes  of  phosphorus  tri- 
chloride (PCls)  vapor. 

PROBLEMS 

1.  1,000,000  molecules  of  hydrogen  will  unite  with  how  many  mole- 
cules of  oxygen  to  form  how  many  molecules  of  water  vapor?    What  will 
be  the  relative  weights  of  hydrogen  and  water  vapor? 

2.  A  liter  of  sulphurous  oxide  gas  (SO2)  weighs  2.8672  gm.    What  is  the 
molecular  weight  of  this  compound? 

3.  Sodium  chloride  contains  39.32  per  cent  of  sodium,  and  its  molecular 
weight  is  58.5.    What  is  the  atomic  weight  of  sodium,  if  there  is  but  one  atom 
of  sodium  in  a  molecule  of  the  salt? 

4.  Calculate  the  atomic  weight  of  phosphorus  from  the  following  data: 
In  100  parts  of  phosphoric  oxide  there  are  43.66  parts  of  phosphorus  and  56.34 
of  oxygen.    Each  molecule  contains  seven  atoms,  two  of  phosphorus  and 
five  of  oxygen. 

6.  (a)  By  a  certain  method  the  atomic  weight  of  copper  might  be  either 
63.1  or  126.2.  The  specific  heat  of  copper  is  .093.  Which  value  is  prob- 
ably the  atomic  weight  of  copper?  (b)  Similarly,  the  values  for  arsenic  are 
74.8  and  149.6.  The  specific  heat  of  arsenic  is  .082.  Which  weight  is 
probably  correct? 

6.  Determine  the  atomic  weight  of  nitrogen  from  the  following  analyses: 
(a)  A  compound  having  a  molecular  weight  of  65  is  21.539  per  cent  nitro- 
gen.   (6)  Similarly  101  and  13.861.    (c)  63  and  22.222.  (d)  44  and  63.63. 

7.  If  2  gm.  of  potassium  chloride  yield  3.84  gm.  of  silver  chloride,  cal- 
culate the  atomic  weight  of  potassium. 


CHAPTER   XIII 
VALENCE  — EQUIVALENT  WEIGHT 

145.  Valence  of  Atoms  and  Groups  of  Atoms. —  Formu- 
las obtained  by  the  methods  discussed  in  Chapter  XII 
show  certain  regularities.  Let  us  take  as  examples  the 
following  groups :  - 

I.   HYDROGEN  COMPOUNDS 

CH4 
SiH4 


P205          S03 
CrO3 


HC1 
HBr 

H2O 
H2S 

NH3 
PH3 

II.   OXIDES 

Na2O 
K2O 

CuO 

MgO 

A1203 
P203 

S02 
C02 

III.   ACIDS  AND  SALTS 


HC1 

HNO3 

H2S04 

H3P04 

NaCl 

NaNO3 

Na2SO4 

Na3P04 

CaCl2 

Ca(N03)2 

CaSO4 

Ca3(P04)2 

AlCla 

A1(N03)3 

A12(S04)3 

A12(P04)3 

IV.   BASES 

NaOH         Ca(OH)2        A1(OH)3 
KOH          Ba(OH)2         Bi(OH)3 

These  groups  might  be  greatly  extended.  Careful  com- 
parison of  these  and  many  other  formulas  leads  to  two 
conclusions,  (i)  Atoms  of  elements  differ  in  the  number 


VALENCE  127 

of  atoms  or  atomic  groups  of  other  elements  with  which 
they  combine.  Thus,  in  Group  I  one  atom  each  of  chlorine 
and  bromine  combines  with  one  of  hydrogen;  one  atom  of 
oxygen  and  of  sulphur  combines  with  two  of  hydrogen; 
and  so  on.  (2)  One  atom  of  certain  elements  unites  with 
only  one  atom  or  atomic  group  of  certain  other  elements, 
with  only  two  atoms  or  two  atomic  groups  of  certain 
others,  and  so  on.  These  two  conclusions  mean  that 
atoms  of  elements  have  a  definite  combining  power. 
The  number  which  expresses  the  combining  power  of  an 
atom  of  an  element  is  called  the  valence  of  the  element. 

Strictly  speaking,  valence  is  a  property  of  atoms. 
But  many  atomic  groups  often  act  chemically  like  indi- 
vidual atoms,  e.g.  they  pass  as  a  whole  from  compound  to 
compound ;  hence  it  is  customary  to  speak  of  the  valence 
of  atomic  groups  like  SO4,  NO3,  NH4,  and  OH.  Such 
groups  are  often  called  radicals  (105);  occasionally  the 
term  radical  is  applied  by  analogy  to  a  single  atom,  e.g. 
Cl  in  chlorides,  C  in  carbides,  S  in  sulphides  (as  in 
Table  B,  page  128). 

146.  Tables  of  Valence.  —  The  valence  of  certain  ele- 
ments and  radicals  is  shown  in  Tables  A  and  B.  These 
tables  are  very  useful  in  writing  formulas  (149). 

147.  Valence  Terms.  —  An  element  or  radical  which 
has  the  valence  I  is  called  a  monad;  those  which  have 
the  valence  II  are  called  dyads;  similarly,  those  elements 
or  groups  which  have  the  valence  III,  IV,  V,  VI  are 
called  respectively  triads,  tetrads,  pentads,  and  hexads. 
The  corresponding  adjectives  are  univalent,  bivalent  (or 
divalent),  trivalent,  quadrivalent  (or  tetravalent) ,  quin- 
quivalent (or  pentavalent),  and  hexavalent.  The  valence 
of  an  element  or  radical  is  often  represented  by  small 
Roman  numerals,  e.g.  H1,  On,  Alm,  SO4n,  OH1. 


128  CHEMISTRY 

TABLE  A  —  VALENCE  OF  CERTAIN  ELEMENTS 


Name 

Symbol 

Valence 

Name 

Symbol 

Valence 

Aluminium    

Al 
(NH4) 
Sb 
Sb 
As 
As 
Ba 
Bi 
Ca 
C 
Cr 
Cu 
Cu 
Au 
Au 
H 

III 
I 
III 
V 
III 
V 
II 
III 
II 
IV 
III 
I 
II 
I 
III 
I 

Iron  (ous)  . 

Fe 
Fe 
Pb 
Mg 
Hg 
Hg 
P 
P 
K 
Ra 
Si 
Ag 
Na 
Sn 
Sn 
Zn 

II 
III 
II 
II 
I 
II 
III 
V 
I 
II 
IV 
I 
I 
II 
IV 

II 

(Ammonium)    
Antimony  (ous)  .... 
Antimony  (ic)  ..... 
Arsenic  (ous) 

Iron  (ic)   
Lead  
Magnesium         .... 

Mercury  (ous) 

Arsenic  (ic)    

Mercury  (ic)  
Phosphorus  (ous)  .  .  . 
Phosphorus  (ic)  .... 
Potassium  
Radium  

Barium  
Bismuth 

Calcium       

Carbon 

Chromium 

Silicon          

Copper  (ous)  
Copper  (ic)  
Gold  (ous)  

Silver  

Sodium 

Tin  (ous)  
Tin  (ic) 

Gold  (ic) 

Hydrogen  

Zinc  

TABLE  B  —  VALENCE  OF  CERTAIN  RADICALS 


Group  Name  of 
Compound 

Symbol 
or 
Formula 

Valence 

Group  Name  of 
Compound 

Symbol 
or 
Formula 

Valence 

Acetate 

C2H3O2 

I 

Iodide  

I 

I 

Bromide  
Carbide 

Br 

c 

I 
IV 

Manganate  
Nitrate 

MnO4 
NO3 

II 
I 

Carbonate 

CO3 

II 

Nitrite      

NO2 

I 

Carbonate  (acid)  . 
Chlorate 

HC03 
C1O3 

I 
I 

Oxide  

Permanganate.  .  .  . 

0 
Mn04 

II 

I 

Chloride  
Chromate 

Cl 
CrO4 

I 
II 

Phosphate  (ortho) 
Silicate  (meta)  .  .  . 

P04 
Si03 

III 
II 

Dichromate 

Cr2O7 

II 

Sulphate           .... 

SO4 

II 

Ferricyanide  .... 
Ferrocyanide 

Fe(CN)6 
Fe(CN)6 

III 
IV 

Sulphate  (acid)  .  .  . 
Sulphide       

HSO4 

s 

I 
II 

Fluoride  
Hydroxide 

F 
OH 

I 
I 

Sulphite  
Sulphite  (acid)  .  .  . 

S03 
HS03 

II 

I 

VALENCE  129 

The  valence  of  some  elements  varies  with  the  combining 
element  and  the  conditions  under  which  the  combina- 
tion occurred.  The  valence  of  most  radicals  is  unvaried. 
For  convenience  in  learning  the  valence  of  radicals  and 
the  usual  valence  of  some  elements,  the  valences  given 
in  Tables  A  and  B  may  be  rearranged  as  follows :  — 

I.  Monads  — Ag,    Au(ous),    Br(ide),     C2H3O2,    Cl(ide),    C1O3, 
Cu(ous),  F(ide),  H,  HCO3,  HSO3,  HSO4,  Hg(ous),  I,  K,  MnO4(per), 
Na,  NH4,  N02,  NO3,  OH. 

II.  Dyads  — Ba,  Ca,  CO3,  CrO4,  Cr2O7,  Cu(ic),  Fe(ous),  Hg(ic), 
Mg,  MnO4(ate),  O,  Pb,  Ra,  S(ide),  S03,  SO4,  SiO3,  Sn(ous),  Zn. 

III.  Triads  —  Al,  As(ous),  Au(ic),  Bi,   Cr,   Fe(ic),   Fe(CN)6(i), 
P(ous),  P04(o),  Sb(ous). 

IV.  Tetrads  — C,  C(ide),  Fe(CN).(o),  Si,  Sn(ic). 

V.  Pentads  — As(ic),  P(ic),  Sb(ic). 

148.  Valence  illustrated  by  Combination  and  Sub- 
stitution. —  Elements  and  atomic  groups  which  have  the 
same  valence  combine  with  each  other  unit  for  unit. 
Thus,  one  atom  of  sodium  combines  with  one  atom  of 
chlorine,  to  form  one  molecule  of  sodium  chloride;  and 
one  NH4-group  combines  with  one  atom  of  chlorine  to 
form  one  molecule  of  ammonium  chloride.  Elements 
and  atomic  groups  having  a  different  valence  usually 
(151)  combine  with  each  other  so  that  the  total  valence 
of  the  atoms  of  each  element  or  group  in  a  molecule  is 
equal.  Thus,  two  atoms  of  hydrogen  (each  having  the 
valence  I)  combine  with  one  of  oxygen  (having  the  valence 
II)  to  form  one  molecule  of  water  (H2O) ;  and  two  atoms 
of  aluminium  (each  having  the  valence  III)  combine 
with  three  SO4-groups  (each  having  the  valence  II)  to 
form  one  molecule  of  aluminium  sulphate  (A12(SO4)3). 

Just  as  atoms  and  groups  of  the  same  valence  combine 
unit  for  unit,  or  those  of  different  valence  combine  so 


130  CHEMISTRY 

that  the  valences  ef  each  element  or  atomic  group  bal- 
ance, so  also  atoms  and  atomic  groups  displace  each 
other  —  unit  for  unit  if  the  valence  is  the  same,  or  equiva- 
lently  if  the  valence  is  different.  For  example,  when  zinc 
is  substituted  for  hydrogen  in  hydrochloric  acid,  one 
atom  of  zinc  (having  the  valence  II)  displaces  two  atoms 
of  hydrogen  (each  having  the  valence  I),  and  the  formula 
of  the  resulting  zinc  chloride  is  ZnCl2 ;  similarly,  one  atom 
of  chlorine  (having  the  valence  I)  displaces  one  OH-group 
from  sodium  hydroxide,  and  the  formula  of  the  resulting 
sodium  chloride  is  NaCl. 

149.  Writing  Formulas  from  Valence.  —  Formulas  of 
many  compounds  may  be  written  by  using  the  tables  of 
valence.     Suppose  the  formula  of  magnesium  chloride  is 
desired.    From  Table  A  the  valence  of  magnesium  is  found 
to  be  II,  and  from  Table  B  the  valence  of  chlorine  in 
chlorides  is  found  to  be  I.    Remembering  the  rule  that  in 
most  compounds  the  total  valence  of  the  two  parts  must 
be  equal,  it  is  necessary  to  have  two  atoms  of  chlorine  and 
one  of  magnesium;   hence  the  formula  is  MgCl2.     Again, 
suppose  we  wish  the  formula  of  lead  nitrate.     From  the 
tables  the  valence  of  lead  is  II  and  of   the  NOs-group 
is  I;    therefore  it  is  necessary  to1  have  two  NO3-groups 
for  one  atom    of   lead,   and    the   formula   is   Pb(NO3)2- 
Similarly,    the   formula    of    aluminium    oxide   is   A12O3, 
because  2  and  3  are  the  least  number  of  atoms  of  Al  and 
O  which  give  equal  valence  (VI  in  each  case).    Formulas 
of  acids  may  be  written  by  prefixing  the  correct  number  of 
hydrogen  atoms  to  the  atom  or  group  selected  from  Table 
B.    Thus,  HNO2  is  the  formula  of  nitrous  acid,  and  HBr 
of  hydrogen  bromide.     (See  Exercises  at  the  end  of  this 
chapter.) 

150.  Multiple  Valence.  —  Certain  elements  have  two 


VALENCE  131 

values  for  their  valence  (Tables  A  and  B,  and  147). 
Thus,  copper  is  I  or  II,  iron  II  or  III,  phosphorus  III  or 
V;  the  valence  in  each  case  depends  upon  the  kind  of 
compound  of  which  the  element  is  a  constituent,  i.e. 
whether  cuprous  or  cupric,  ferrous  or  ferric,  phosphorous 
or  phosphoric.  Some  elements  have  several  values  for 
their  valence.  Thus,  sulphur  is  II  in  hydrogen  sulphide 
(H2S)  and  all  other  sulphides,  IV  in  sulphur  dioxide 
(SO2),  and  VI  in  sulphur  trioxide  (SO3);  nitrogen  is  I  in 
nitrous  oxide  (N2O)  ,  II  in  nitric  oxide  (NO)  ,  IV  in  nitrogen 
dioxide  (NO2),  and  V  in  nitrogen  pentoxide  (N2O5). 

151.  Exceptional  Compounds.  —  The  formulas  of  cer- 
tain compounds  apparently  do  not  conform  to  the  sim- 
plified conception  of  valence  discussed  above.     Carbon 
monoxide  (CO),  calcium  carbide  (CaC2),  lead  tetroxide 
(Pb3O4),  magnetic  iron  oxide  (Fe3O4),  hydrogen  dioxide 
(H2O2),  acetylene  (C2H2),  and  ethylene  (C2H4)   are  the 
common   exceptions.     Interpretation   of   these   apparent 
exceptions  must  be  sought  in  a  larger  text  book. 

152.  Representation    of    Valence.  —  The    valence    of 
elements    and    atomic    groups    may    be    represented    in 
various  ways.     One    has  already  been  given,   viz.   H1, 
On,  PO4m.    Sometimes  short  lines  are  used,  e.g.  H  -  ,  O  =  , 
—  0  —  ,  Al  =  ,  etc.    If  lines  are  used  to  represent  valence  in 
compounds,  a  single  line  answers  for  two  elements.    Thus, 
the  formula  of  water  is  written  H-O-H  rather  than 

/H 

H  --  O  --  H.     Similarly,   ammonia   gas   is   N—  H  and 


XOH  /0-H 

calcium   hydroxide   is   Cav  or    Cax  •       Such 

XOH  XO-H 


132  CHEMISTRY 

formulas  are  called  structural  or  graphic  formulas,  for 
they  show  the  probable  arrangement  of  the  atoms  in 
molecules.  Thus,  the  graphic  formula  of  nitric  acid  is 

O 
usually  written  H  —  O  —  N^     ,  because  this  arrangement 

xo 

of  atoms  not  only  displays  the  correct  valence  of  each 
element  but  it  also  represents  certain  facts  about  nitric 
acid,  e.g.  that  the  hydrogen  is  not  combined  directly 
with  nitrogen.  Structural  formulas  are  useful,  especially 
inorganic  chemistry  (Chapter  XVII),  but  it  must  not  be 
forgotten  that  they  are  merely  representations;  the  lines 
are  intended  to  indicate  the  numerical  value  of  the  valence 
and  not  the  strength  of  the  combination  of  the  atoms. 

Equivalent  Weight 

153.  Determination  of  Valence  from  Atomic  and 
Equivalent  Weights.  —  The  valence  of  an  element  is 
found  by  dividing  its  atomic  weight  by  its  equivalent 
weight,  that  is :  - 

Valence  =  Atomic  Weight  -f-  Equivalent  Weight. 

The  term  equivalent  weight  requires  explanation.  Equiva- 
lent weights  of  the  elements  are  the  weights  that  are 
chemically  equivalent,  provided  the  weights  are  expressed 
in  terms  of  the  standard,  viz.  8  grams  of  oxygen.  In  other 
words,  the  equivalent  weight  of  an  element  is  the  weight 
that  combines  with  or  replaces  8  grams  of  oxygen.  For 
example,  it  is  found  by  experiment  that  magnesium  and 
oxygen  combine  in  the  ratio  of  60  to  40  respectively  (in 
round  numbers)..  Now  if  for  the  40  we  substitute  8, 
then  the  60  becomes  12,  that  is,  12  is  the  number  of 


EQUIVALENT    WEIGHTS 


grams  of  magnesium  that  unites  with  8  grams  of  oxygen. 
Therefore  12  is  the  equivalent  weight  of  magnesium. 
Equivalent  weights  of  the  elements  have  been  carefully 
determined  by  experiment.  When  the  atomic  weights 
are  divided  by  the  equivalent  weights,  the  quotient  is 
the  valence  of  the  respective  elements  (Table  C). 

TABLE    C  —  EQUIVALENT    WEIGHTS,    ATOMIC    WEIGHTS,    AND    VALENCE 


Element 

Equivalent 
Weight 

Atomic 
Weight 

Valence 

Oxygen                                              .    . 

8  oo 

16  oo 

2 

Aluminium  
Bromine                                       

9-03 
70.02 

27.1 

70  02 

3 
i 

Calcium 

20  03  5 

4O  O7 

2 

Carbon  ....                     

3.00 

I2.OO 

4" 

Chlorine                                          .... 

35.46 

?r  46 

I 

Copper  (ous)  

63.57 

6^.57 

I 

Copper  (ic)  .  .               

31.785 

63.  <7 

2 

Hydrogen                                        .... 

i.  008 

I  008 

I 

Iron  (ous)  

27.92 

55.84 

2 

Iron  (ie)                                         

18.62 

55.84 

•} 

Magnesium                                         .  . 

12.  l6 

24.72 

2 

Mercury  (ous)            

2OO.6 

2OO.6 

I 

Mercury  (ic)                                 

100.3 

2OO.6 

2 

Potassium                                             •  . 

30.  IO 

30.  IO 

I 

Silver.             

107.88 

107.88 

I 

Sodium                                 

23.00 

23.OO 

I 

Sulphur                                           .    •  • 

16.035 

32.O7 

2 

Zinc 

32.685 

65.37 

2 

154.   Equivalent  Weights  and  their  Determination.  - 

The  list  of  equivalent  weights  given  above  might  be 
extended  to  include  all  the  elements  which  form  com- 
pounds. These  numbers  are  sometimes  called  combining 
numbers,  combining  weights,  or  simply  equivalents. 
The  term  equivalent  weights  is  preferable,  because  they 


134  CHEMISTRY 

actually  are  the  weights  chemically  equivalent  to  each 
other.  Thus,  if  we  start  with  hydrogen  chloride  (HC1), 
1.008  gm.  of  hydrogen  —  to  take  a  convenient  denomina- 
tion —  combines  with  35.46  gm.  of  chlorine,  and  this 
1.008  gm.  of  hydrogen  can  be  replaced  chemically  by  32.685 
gm.  of  zinc,  12.16  gm.  of  magnesium,  39.10  gm.  of  potas- 
sium, or  23.00  gm.  of  sodium,  and  so  on.  These  elements 
are  chemically  equivalent  in  the  ratio  of  these  weights. 

Equivalent  weights  can  be  readily  found  by  experiment,  though 
the  equivalent  weight  is  not  always  found  directly  in  terms  of 
oxygen.  The  equivalent  weight  of  magnesium  can  be  found  by 
heating  a  known  weight  of  magnesium  in  the  air,  taking  care,  of 
course,  not  to  lose  the  product.  The  interaction  of  metals  and 
acids  provides  a  simple  method  of  finding  the  equivalent  weights 
of  zinc,  magnesium,  aluminium,  and  iron;  the  interaction  of  water 
with  sodium  and  with  calcium  permits  the  determination  of  the 
equivalent  weights  of  these  metals;  and  the  displacement  of  metals 
from  solutions  of  their  compounds  by  zinc  provides  another  general 
method.  Illustrations:  Suppose  .4188  gm.  of  iron  liberated  186  cc. 
of  hydrogen  at  20°  C.  and  754.5  mm.  The  volume  of  hydrogen  is 
168  cc.,  when  reduced  to  standard  conditions  (Chapter  IV);  its  weight 
is  .015  gm.  (i.e.  .168  X  .0898).  Therefore  .4188  :  .015  ::  x  :  i,  x  = 
27.92,  which  is  the  equivalent  of  iron  (ous).  Again  suppose  4.085  gm. 
of  zinc  combine  with  i  gm.  of  oxygen  to  form  zinc  oxide.  The 
proportion  is  4.085  :  i  ::  x  :  8,  or  x  =  32.68,  which  is  the  equivalent 
of  zinc. 

EXERCISES 

1.  Select  three  formulas  and  illustrate  valence  by  them. 

2.  State  the  fundamental  rule  for  writing  formulas  from  valence. 

3.  Write  the  formula  of    the  chloride  of   potassium,  sodium,  silver, 
copper  (ous),  copper  (ic),  mercury  (ous),  mercury  (ic),  iron  (ous),  iron  (ic), 
zinc,  tin  (ous),  tin  (ic),  calcium,  barium,  magnesium,  bismuth,  aluminium, 
ammonium,  lead.    (Use  Valence  Tables.) 

4.  Write  the  formula  of  the  sulphate  of  K,  Na,  Ag,  Cu,  Fe  (ous),  Zn, 
Pb,  Ca,  Ba,  Mg,  Cr,  Al,  NH4.     (Use  Valence  Tables.) 

5.  Write  the  formulas  of  the  following  compounds  and  indicate  the 
valence  of  each  element  or  radical  by  Roman  numerals :  ammonium  fluoride, 


VALENCE  135 

sodium  silicate,  potassium  manganate,  barium  phosphate,  zinc  iodide, 
ammonium  chromate,  silver  chromate,  magnesium  oxide,  sodium  dichro- 
mate,  aluminium  chloride,  ferrous  bromide,  calcium  phosphate,  mercurous 
nitrate. 

6.  Write  the  formulas  of  the  following  (as  in  Exercise  12):   ferrous 
carbonate,  aluminium  phosphate,  calcium   fluoride,  sodium  permanganate, 
phosphoric  acid,  silicic  acid,  sulphurous  acid,  nitrous  acid,  chromic  acid, 
hydriodic  acid,  carbonic  acid. 

7.  Define  and  illustrate    the  term    equivalent  weight.      What  is  the 
equivalent  weight  of  hydrogen,  oxygen,  sulphur,  zinc,  copper,  magnesium, 
silver,  potassium,  aluminium? 

8.  Write  the  formula  of  the  nitrate  of  calcium,  silver,  iron  (ic),  mer- 
cury (ous),  barium,  magnesium,  and  aluminium.     (Use  Valence  Tables.) 

9.  Write  the  formula  of  the  oxide  of  Al,  Sb  (ous  and  ic),  As  (ous  and  ic), 
Ba,  Bi,  Ca,  Cr.     (Use  Valence  Tables.) 

10.  Write  the  formula  of  (a)  the  carbonate  of  Cu  (ic),  Ba,  Ag;    (ft)  the 
iodide  of  Hg  (ous  and  ic),  Pb,  Ag,  Na,  Fe  (ous  and  ic),  Al;  (c)  the  chlorate 
of  Na,  Ba,  Ag.     (Use  Valence  Tables.) 

11.  Write  the  formula  of  (a)  the  dichromate  of  Pb,  Na,  Al,  Ag;   (ft)  the 
manganate  of  Na,  Ca,  Ag;    (c)  the  nitrate  of  Pb,  Ca,  Cu  (ous  and  ic).    (Use 
Valence  Tables.) 

12.  Write  the  formula  of  (a)  the  nitrite  of  Na,  Ca,  Pb,  Cu  (ic);    (ft) 
the  permanganate  of  Na,  Ag,  H,  K;   (c)  the  sulphate,  chloride,  nitrate,  car- 
bonate, sulphite,  bromide  of  H.     (Use  Valence  Tables.) 

PROBLEMS 

1.  Calculate  the  equivalent  weights  of  the  respective  metals  from  the 
following  data:   (a)  .5  gm.  of  calcium  unites  with  .2  gm.  of  oxygen  to  form 
calcium  oxide  (CaO).      (ft)  15  gm.  of  mercury  unite  with  1.2  gm.  of  oxygen 
to  form  mercuric  oxide  (HgO). 

2.  Calculate  the  equivalent  weight  of  sodium  from  the  following:    (a) 
2.3  gm.  of  sodium  liberate  .1  gm.  of  hydrogen  from  water,     (ft)   1.15  gm.  of 
sodium  liberate  555.5  cc.  of  hydrogen  (at  standard  conditions). 

3.  If  .03  gm.  of  magnesium  yields  30.4  cc.  of  hydrogen  at  20°  C.  and  750 
mm.,  what  is  the  equivalent  weight  of  magnesium? 

4.  If  45.731  gm.  of  silver  form  60.749  8m-  of  silver  chloride  (AgCl), 
calculate  the  equivalent  weight  of  silver.     (Assume  35.46  as  the  equivalent 
weight  of  chlorine.) 

5.  Calculate  the  equivalent  weight  of  aluminium  from  the  following: 
(a)  .17  gm.  of  aluminium  liberates  221  cc.  of  hydrogen  at  10°  C.  and  754 
mm.    (ft)  Weight  of  aluminium  =  .163  gm.,  t  =  13.5°  C.,  pres.  =  754  mm., 
vol.  of  hydrogen  =  217  cc. 


CHAPTER  XIV 


SOLUTION  —  ACIDS,  BASES,  AND   SALTS 

155.  Introduction.  —  Many  properties  of  solutions  and 
the  general  properties  of  acids,  bases,  and  salts  have  already 
been  studied  (42-51,  90-94) ;   considerable  attention  has 
also  been  given  to  certain  examples  of  acids,  bases,  and 
salts,  viz.  hydrochloric  and  nitric  acids,  ammonium  hy- 
droxide,  chlorides    and   nitrates  (82-89,   105-113).     In 
the  present  chapter  we  shall  discuss  more  fully  the  prop- 
erties  of   acids,   bases,  and  salts,  and  continue  the   con- 
sideration of  solutions,  especially  the  nature  of  solutions 
of  acids,  bases,  and  salts. 

156.  Behavior  of  Solutions  toward  an  Electric  Current. 
-  Water  itself  conducts  electricity  very  slightly  indeed. 

Therefore,  if  substances 
are  dissolved  in  water 
and  the  solutions  are 
subjected  to  the  action 
of  an  electric  current, 
it  can  be  readily  found 
what  substances  are 
electrolytes  and  what 
are  non-electrolytes,  i.e. 
what  substances  form 
conducting  solutions  and 
what  form  non-conducting.  The  apparatus  shown  in 
Fig.  32  may  be  used  for  this  experiment.  It  consists  of 
a  battery  (A)  and  an  electric  bell  (B)  connected  with 


Fig.  32.  —  Apparatus  for  Showing  the 
Behavior  of  Solutions  toward  an 
Electric  Current. 


ACIDS,   BASES,   AND   SALTS  137 

pieces  of  carbon  (C,  D),  called  electrodes,  through  which 
the  electric  current  enters  and  leaves  the  solution  in  the 
vessel  (E).  One  electrode  is  called  the  positive  (+)  elec- 
trode or  anode,  and  the  other  the  negative  (— )  electrode 
or  cathode.  When  a  current  enters  the  system  (battery, 
bell,  solution),  the  bell  rings  only  when  the  solution  con- 
tains an  electrolyte.  Experiment  shows  that  only  acids, 
bases,  and  salts  are  electrolytes.  Among  non-electrolytes 
are  sugar,  alcohol,  chloroform,  and  glycerin.  (See  Part 
II,  Exp.  72.) 

157.  Theory  of  Electrolytic  Dissociation.  —  A  theory 
was  proposed  in  1887  by  the  Swedish  physicist  Arrhenius 
to  explain  the  properties  of  dilute  solutions.    It  is  called 
the  theory  of  electrolytic  dissociation  and  may  be  stated 
as  follows :  — 

When  acids,  bases,  and  salts  are  dissolved  in  water,  their 
molecules  dissociate  to  a  varying  degree  into  independent 
particles  charged  with  electricity. 

This  theory  means  that  a  solution  of  sodium  chloride, 
for  example,  consists  of  water  throughout  which  are  dis- 
tributed some  molecules  of  sodium  chloride  together 
with  electrically-charged  particles  into  which  the  remain- 
ing molecules  of  sodium  chloride  have  dissociated. 

158.  The  Theory  Expanded  and  Terms  Defined.  —  The 
dissociation  of  acids,  bases,  and  salts  when  in  aqueous 
solution  is  called  electrolytic  dissociation  or  ionization. 
The  independent  particles  that  are  charged  with  electricity 
are  called  ions.     Each  ion  is  a  portion  of  a  molecule  and, 
as  we  have  said,  is  charged  with  electricity.    Two  kinds 
of  ions  are  present  in  every  electrolytic  solution,  viz.  elec- 
tro-positive ions,   or  cations,   and  electro-negative  ions, 
or  anions.     Ions,  although  formed  by  the  dissociation  of 
molecules,  are  not  identical  with  atoms,  but  differ  mainly 


138  CHEMISTRY 

in  having  a  charge  of  electricity.  For  example,  in  a  solu- 
tion of  sodium  chloride  the  electro-positive  sodium  ions 
move  about  in  the  water  without  producing  any  apparent 
chemical  change;  but  ordinary  sodium  interacts  vio- 
lently with  water,  as  we  have  already  seen  (21).  Simi- 
larly, the  chloride  ions  circulate  freely  in  water  and 
exhibit  none  of  the  effects  of  gaseous  chlorine  on  water 
(53) .  Ions  may  be  re-defined  as  electrically  charged  atoms 
or  atomic  groups.  It  should  not  be  overlooked  that  the 
electric  charges  are  on  ions  into  which  molecules  have 
dissociated,  and  are  not  acquired  from  the  electricity 
that  may  subsequently  be  passed  through  the  solution. 

159.  Representation  of  Ions.  —  It  is  customary  to  rep- 
resent ions  by  chemical  symbols  supplemented  by  the  sign 
that  designates   the  kind  of   electric  charge.     Thus,  the 
ions  formed  by  the  dissociation  of  sodium  chloride  are 
Na+  and  Cl~,  and  of  potassium  nitrate  are  K+  and  NOs~. 
In  sodium  chloride  solution,  the  number  of  sodium  ions 
equals  the  number  of  chloride  ions  and  the  sum  of  the 
positive  charges  on  the  sodium  ions  equals  the  sum  of 
the  negative  charges  on  the  chloride  ions.     In  the  case 
of  calcium  chloride  (CaCl2),  each  molecule  dissociates  into 
two  chloride  ions  and  one  calcium  ion;  but  since  the  sum 
of  each  kind  of  electric  charges  must  be  the  same,  each 
calcium  ion  must  carry   twice   the  charge  which  is  on 
each  chloride  ion.     Hence  the  ions  formed  by  the  disso- 
ciation of  calcium  chloride  are  designated  Ca++  and  2C1~. 
(See  Table  of  lonization,  168.) 

160.  lonization  and  Acids,  Bases,  and  Salts.  —  As  pre- 
viously stated,  only  acids,  bases,  and  salts  are  electrolytes. 
The  general  properties  of  these  substances  have  already 
been  discussed  (90-94)  and  their  special  characteristics 
are  treated  under  the  individual  compounds.     (See  also 


ACIDS,   BASES,   AND   SALTS  139 

164-168.)  We  can  now  regard  them  as  a  single  class 
of  substances  (i.e.  electrolytes)  which  dissociate  in  solu- 
tion into  ions.  For  example,  a  solution  of  hydrochloric 
acid  contains  hydrogen  ions  and  chloride  ions,  while  in 
sodium  hydroxide  solution  there  are  sodium  ions  and 
hydroxyl  ions,  and  in  silver  nitrate  solution  silver  ions 
and  nitrate  ions. 

161.  Interpretation  of  Certain  Facts  by  the  Theory  of 
Electrolytic    Dissociation.  —  The    theory    of    electrolytic 
dissociation,  like  other  theories,  is  of  value  only  as  far 
as  it  is  an  acceptable  explanation  of  facts  derived  from 
experiment.    We  shall  now  interpret  certain  facts  by  the 
theory  of  electrolytic  dissociation,  especially  facts  which 
have  been  observed  in  connection  with  solutions  of  acids, 
bases,  and  salts. 

162.  Freezing  Points  and  Boiling  Points  of  Solutions.— 
It  has  long  been  known  that  solutions  boil  at  a  higher 
temperature  and  freeze  at  a  lower  temperature  than  pure 
water.     Now,  when  weights  of  non-electrolytes  equal  to 
their  molecular  weights  are  dissolved  in  the  same  weight 
of  water,   experiment  shows   that   the  boiling  point  of 
each  solution  is  raised  approximately  the  same  number 
of   degrees;  analogous  statements  can  be  made  about  the 
freezing  point.     Thus,  the  boiling  point  of  water  is  raised 
about  .52°  C.  and  the  freezing  point  lowered  about  1.89°  C. 
by  a  solution  of  46  gm.  of  alcohol  (C2H60)  in  1000  gm.  of 
water.    But  when  solutions  of  electrolytes,  i.e.  acids,  bases, 
and  salts,  are  experimented  with,  the  freezing  point  is 
lower   than    that    produced    by   non-electrolytes    under 
parallel  conditions.     For  example,  the  addition  of   58.5 
gm.  of  sodium  chloride  to  1000  gm.  of  water  lowers  the 
freezing  point  of  the  water  about  3.5°C.;    that  is,  the 
depression  of  the  freezing  point  is  nearly  twice  that  pro- 


140  CHEMISTRY 

duced  by  an  alcohol  solution  of  equivalent  concentra- 
tion. These  facts  are  in  harmony  with  the  theory  of 
electrolytic  dissociation.  According  to  the  theory,  solu- 
tions of  non-electrolytes  contain  only  molecules,  while 
solutions  of  electrolytes  contain  molecules  and  ions  — 
the  number  depending  on  the  substance  and  the  concen- 
tration of  the  solution.  Hence  the  number  of  inde- 
pendent particles  (molecules  and  ions)  in  the  electrolytic 
solution  is  greater  than  in  th,e  non-electrolytic  solution. 
Since  ions  have  the  same  effect  as  molecules  on  the 
boiling  and  freezing  points  of  a  solution,  the  larger  the 
number  of  particles  the  greater  should  be  the  raising  of 
the  boiling  point  and  the  lowering  of  the  freezing  point. 
Moreover,  when  molecular  weights  are  determined  by 
the  freezing  point  or  the  boiling  point  method,  the  results 
still  further  confirm  the  theory.  Thus,  (a)  the  molec- 
ular weights  of  non-electrolytes  agree  with  the  values 
found  by  other  methods,  while  (b)  the  molecular  weights 
of  electrolytes  (i.e.  acids,  bases,  and  salts)  are  abnormal, 
that  is,  not  only  do  the  values  of  electrolytes  disagree 
with  those  found  by  other  methods  but  the  values  also 
vary  with  the  concentration  of  the  solution.  Therefore, 
as  far  as  freezing  and  boiling  points  are  concerned,  the 
theory  furnishes  an  acceptable  interpretation  of  the 
facts. 

163.  Chemical  Behavior  of  Electrolytes  in  Solution.  - 
Dry  potassium  chloride  and  dry  silver  nitrate  do  not 
interact  chemically,  but  if  their  solutions  are  mixed,  a 
precipitate  of  silver  chloride  is  immediately  produced. 
Furthermore,  any  dissolved  chloride  will  interact  in  the 
same  way  with  silver  nitrate.  Let  us  interpret  these  two 
sets  of  facts  by  the  theory  of  electrolytic  dissociation. 
Solutions  of  potassium  chloride  and  silver  nitrate  contain 


ACIDS,   BASES,   AND   SALTS  141 

potassium  ions  (K+),  chloride  ions  (Cl~),  silver  ions 
(Ag+),  and  nitrate  ions  (NO3~).  Now  when  certain  ions 
are  introduced  into  the  same  solution,  they  seek  each 
other  out,  so  to  speak,  and  combine  chemically  in  accord- 
ance with  certain  fundamental  principles  of  chemical 
action.  Thus,  silver  ions  and  chloride  ions,  if  brought  into 
the  same  solution,  always  combine  to  form  insoluble  silver 
chloride,  which  serves  as  evidence  of  the  chemical  change. 
Obviously,  the  source  of  the  silver  ions  and  the  chloride 
ions  is  immaterial.  Hence,  hydrochloric  acid  and  solu- 
tions of  chlorides  precipitate  silver  chloride  from  a  solu- 
tion containing  silver  ions,  whether  the  silver  ions  are 
furnished  by  silver  nitrate,  silver  sulphate,  or  any  other 
soluble  silver  compound.  This  is  the  explanation  offered 
by  the  theory  of  electrolytic  dissociation  for  the  usual 
test  for  hydrochloric  acid  and  all  soluble  chlorides  (89). 
The  ordinary  equation  for  this  test  is :  — 

HC1  +  AgNO3  =  AgCl  +  HN03 

This  equation  may  be  written  as  an  ionic  equation, 
thus:- 

H+  +  CF  +  Ag+  +  NOr  =  AgCl  +H+  +  N03", 
or  in  a  general  form  as  follows :  - 

CF  +  Ag+  =  AgCl 

On  the  other  hand,  not  all  dissolved  substances  interact, 
even  when  they  are  apparently  quite  similar.  Thus,  no 
chemical  action  is  observed  when  solutions  of  potassium 
chlorate  and  silver  nitrate  are  mixed,  despite  the  fact 
that  the  solutions  contain  ions  and  that  chlorine  is  a 
constituent  of  potassium  chlorate.  This  apparent  con- 
tradiction is  readily  explained  by  the  theory.  Potassium 
chlorate  solution  contains  potassium  ions  (K+)  and 


142  CHEMISTRY 

chlorate  ions  (C1O3~),  and  when  silver  nitrate  solution  is 
added,  the  solution  contains  four  kinds  of  ions,  —  po- 
tassium ions  (K+),  chlorate  ions  (C103~),  silver  ions 
(Ag+),  and  nitrate  ions  (NO3~).  But  compounds  which 
might  be  formed  by  the  various  combinations  of  these 
ions  are  soluble.  Hence  the  ions  remain  uncombined,  for 
the  most  part,  in  the  solution.  Silver  nitrate  is  effective 
in  testing  for  hydrochloric  acid  or  a  soluble  chloride,  but 
not  for  other  compounds  containing  chlorine,  such  as 
potassium  chlorate  (KC1O3).  Strictly  speaking,  the  test 
is  not  for  the  element  chlorine,  but  for  ionic  chlorine 
(i.e.  chloride  ions) ;  and  since  the  solution  of  potassium 
chlorate  contains  no  ionic  chlorine,  the  test  fails  with  this 
compound.  Similarly,  the  test  for  sulphuric  acid  and  all 
soluble  sulphates  is  the  formation  of  insoluble  barium 
sulphate  (BaS04)  upon  the  addition  of  a  solution  of 
barium  chloride  (or  any  other  soluble  barium  compound), 
because  the  sulphuric  acid  and  sulphate  solutions  con- 
tain sulphate  ions  (SO4~~),  which  combine  with  barium 
ions  (Ba++)  furnished  by  the  soluble  barium  compound. 
But  this  test  is  not  applicable  to  other  sulphur  com- 
pounds, such  as  sulphides,  sulphites,  and  thiosulphates, 
because  solutions  of  these  sulphur  compounds  do  not  con- 
tain sulphate  ions.  (See  Part  II,  Exps.  66,  67,  73.) 

164.  General  Properties  of  Acids,  Bases,  and  Salts 
—  Neutralization.  —  Certain  properties  are  exhibited  by 
solutions  of  each  of  these  classes  of  compounds.  Thus,  it 
will  be  recalled,  acids  have  a  sour  taste  and  turn  litmus 
red,  bases  have  a  bitter  taste  and  turn  litmus  blue,  typical 
salts  vary  somewhat  in  taste  and  in  action  on  litmus, 
though  many  are  salty  and  do  not  change  litmus;  also 
acids  and  bases  neutralize  each  other,  that  is,  interact  and 
form  salts  and  water  (93).  When  all  these  properties  are 


ACIDS,   BASES,   AND   SALTS  143 

interpreted  by  the  theory  of  electrolytic  dissociation,  inter- 
esting and  important  conclusions  result.  Solutions  of 
acids,  bases,  and  salts  contain  ions,  and  the  properties 
of  such  solutions  are  ascribed  to  the  ions.  Thus,  in  all 
solutions  of  acids  there  are  hydrogen  ions,  while  all  solu- 
tions of  bases  contain  hydroxyl  ions.  According  to  the 
theory,  then,  an  acid  may  be  defined  as  a  compound 
whose  water  solution  contains  hydrogen  ions  (H+) ,  while 
a  base  is  a  compound  whose  water  solution  contains 
hydroxyl  ions  (OH~).  These  definitions  should  be  com- 
pared with  those  previously  given  (90-93,  160).  (See 
Part  II,  Exp.  68.) 

Neutralization,  interpreted  by  the  ionic  theory,  is  the 
combining  of  hydrogen  and  hydroxyl  ions  to  form  mole- 
cules of  water.  Suppose  solutions  of  hydrochloric  acid 
and  potassium  hydroxide  are  mixed  in  the  proper  pro- 
portions. The  mixture  at  first  contains  ions  of  hydrogen, 
chlorine,  potassium,  and  hydroxyl.  But  the  hydrogen 
and  hydroxyl  ions  immediately  unite  to  form  molecules 
of  water,  because  water  does  not  dissociate  to  any  appre- 
ciable extent.  The  final  solution  is  thus  rendered  neutral 
by  the  removal  of  the  hydrogen  and  the  hydroxyl  ions  — 
the  acid  and  basic  constituents  respectively.  The  ionic 
equation  expressing  the  neutralization  of  potassium 
hydroxide  by  hydrochloric  acid  is :  - 

K+  +  OH~  +  H+  +  cr  =  K+  +  cr  +  H2o 

The  potassium  and  chloride  ions  remain  free  and  uncom- 
bined  until  the  solution  is  evaporated;  as  the  concentra- 
tion increases,  the  ions  unite  until  nothing  remains  except 
the  neutral  salt  potassium  chloride.  Neutralization  is 
often  accurately  accomplished  by  means  of  burettes 
(Fig.  33).  (See  Part  II,  Exp.  78.) 


144 


CHEMISTRY 


Burettes  are  graduated  glass  tubes  and  are  so  marked  that  any 
desired  portion  of  the  contents  can  be  drawn  off  by  the  stop-cock 
at  the  lower  end.  In  using  burettes  for  neutralization,  one  is  filled 
to  the  mark  with  a  solution  of  an  acid  and  the 
other  with  a  solution  of  a  base  —  one  solution 
being  of  known  concentration.  A  measured  por- 
tion, say  15  cubic  centimeters  of  the  base,  is 
drawn  off  into  a  beaker,  a  few  drops  of  litmus 
or  phenol-phthalein  solution  is  added,  and  the 
acid  is  dropped  in  until  the  change  in  color 
(after  thorough  stirring)  shows  that  neutraliza- 
tion has  occurred.  The  volume  of  acid  added 
is  read  accurately.  Knowing  the  concentration 
of  the  base  solution  and  the  volume  of  the  acid 
solution,  we  can  calculate  the  exact  weight  of 
the  acid  needed  for  the  neutralization  of  the 
base. 

Heat  is  liberated  when  these  ions  unite  to 

p-1G         Burettes    ^orm  water.     If  neutralization  is  the  combining 

of  hydrogen  ions  with  hydroxyl  ions,  the  same 
amount  of  heat  should  be  liberated  when  a  given  weight  of  water 
is  formed.  Experiment  shows  that  the  heat  of  neutralization,  as  it 
is  called,  is  the  same  in  all  cases  of  neutralization,  provided  the 
solutions  are  dilute  and  other  thermal  changes  do  not  occur.  It  is 
expressed  in  terms  of  a  unit  called  the  small  calorie,  i.e.  the  quantity 
of  heat  necessary  to  raise  i  gm.  of  water  i°  C.  in  temperature.  And 
when  1 8  gm.  of  water  are  formed  by  the  act  of  neutralization,  13,700 
calories  are  liberated.  A  general  equation  for  a  typical  case  of 
neutralization  might  be  written  — 


OH"        + 

Hydroxyl  Ion 


H+        = 

Hydrogen  Ion 


H2O    -f-     13,700  calories 

Water 


165.  Properties  of  Salts.  —  Salts,  as  previously  denned, 
are  the  compounds  formed  by  substituting  a  metal  for 
the  hydrogen  of  an  acid,  e.g.  hydrochloric  acid  (HC1) 
gives  chlorides  (NaCl,  etc.)  and  nitric  acid  gives  nitrates 
(NaNO3,  etc.)  (94).  Some  acids  contain  more  than  one 


ACIDS,   BASES,   AND   SALTS  145 

replaceable  hydrogen  atom  in  a  molecule,  and  can  form 
more  than  one  class  of  salts.  For  example,  monobasic 
acids  (e.g.  HC1,  HNO3)  contain  one  replaceable  hydrogen 
atom,  dibasic  acids  (e.g.  H2SO4,  H2CO3)  two,  and  tribasic 
acids  (e.g.  H3PO4)  three.  Ordinarily  monobasic  acids 
form  only  one  class  of  salts,  dibasic  acids  form  two  classes, 
and  tribasic  acids  three.  Again,  we  might  regard  salts  as 
formed  from  bases  by  the  substitution  of  a  non-metallic 
element  or  group  for  the  hydroxyl  of  the  base;  e.g.  sodium 
hydroxide  (NaOH)  gives  chlorides  or  nitrates  by  substi- 
tuting Cl  or  NO3  for  OH.  Just  as  a  molecule  of  some 
acids  contains  more  than  one  hydrogen  atom,  so  a  mole- 
cule of  some  bases  contains  more  than  one  hydroxyl 
group;  such  bases  can  form  more  than  one  class  of  salts. 
Bases  are  classified  as  monacid,  diacid,  and  triacid 
bases,  etc.,  according  to  the  number  of  the  replaceable 
hydroxyl  groups  present  in  a  molecule.  Thus,  sodium 
hydroxide  (NaOH)  is  monacid,  calcium  hydroxide 
(Ca(OH)2)  diacid,  and  aluminium  hydroxide  (A1(OH)3) 
triacid.  Salts  may  be  regarded  as  formed  from  acids 
or  from  bases  by  substitution.  Based  on  the  methods 
of  formation  just  described,  there  are  three  classes  of 
salts  —  normal,  acid,  and  basic.  Normal  salts  are  those 
formed  by  substituting  (i)  a  metal  for  all  the  hydrogen 
of  an  acid  or  (2)  a  non-metal  or  non-metallic  group  for 
all  the  hydroxyl  groups  of  a  base.  Acid  salts  result 
when  part  of  the  hydrogen  of  an  acid  is  replaced  by  a 
metal.  Basic  salts  are  produced  by  replacing  some  of 
the  hydroxyl  groups  of  a  base  by  a  non-metal  or  non- 
metallic  group.  The  varying  power  of  replaceability  is 
called  the  basicity  of  an  acid  and  the  acidity  of  a  base. 
The  relations  of  acids,  bases,  and  salts  may  be  represented 
thus:  — 


i46 


CHEMISTRY 


Acid 
H2SO4 

Sulphuric  Acid 


Base 

Zn(OH)2 

Zinc  Hydroxide 


-Normal  Salt' 
Na2SO4 

Sodium  Sulphate 

ZnCl2 

Zinc  Chloride 


Basic  Salt 
Zn(OH)Cl 

Basic  Zinc  Chloride 


Acid  Salt 
HNaS04 

Acid  Sodium  Sulphate 

The  terms  normal,  acid,  and  basic  as  applied  to  salts 
indicate  their  composition,  not  their  general  properties. 
For  when  solutions  of  salts  are  tasted  or  tested  with 
litmus,  the  results  vary.  Normal  salts  may  be  neutral, 
acid,  or  basic  toward  litmus;  acid  salts  may  be  acid  or 
nearly  neutral;  most  basic  salts  are  almost  insoluble  in 
water  and  the  solutions  react  only  faintly  with  litmus. 
The  theory  of  electrolytic  dissociation  offers  an  acceptable 
explanation  of  these  facts.  According  to  the  theory, 
solutions  of  salts,  of  course,  contain  ions.  We  should 
expect  solutions  of  acid  salts  to  contain  hydrogen  ions,  of 
basic  salts  to  contain  hydroxyl  ions,  and  of  normal  salts 
to  contain  neither.  Many  solutions  of  salts  —  acid  and 
basic  as  well  as  neutral  —  behave  toward  litmus  as  we 
should  expect.  Thus,  acid  sodium  sulphate  (HNaSO4) 
is  acid,  basic  bismuth  nitrate  (Bi(OH)2NO3)  is  basic,  and 
sodium  chloride  (NaCl)  is  neutral.  But  certain  salts 
behave  abnormally  toward  litmus.  Thus,  a  solution  of 
the  normal  salt  copper  sulphate  (CuSO4)  is  acid,  of  normal 
sodium  carbonate  is  basic,  and  of  acid  sodium  carbonate 
is  neutral.  Why?  Hitherto  pure  water  has  been  called  a 


ACIDS,   BASES,    AND   SALTS  147 

non-electrolyte,  and  indeed  in  most  cases  there  is  only  the 
very  slightest  evidence  of  its  dissociation  into  the  ions 
H+  and  OH~.  Under  certain  conditions,  however,  the 
slight  ionization  of  water  becomes  a  significant  factor  in 
solutions;  i.e.  its  ions  participate  in  certain  changes. 
Let  us  take  as  examples  the  three  apparently  abnormal 
compounds  just  mentioned,  (i)  Sodium  carbonate  yields 
the  ions  2Na+  and  CO3~~~.  But  the  C03-ions  are  unstable 
and  tend  to  combine  with  H-ions  to  form  HCO3-ions 
(HC03~).  The  gradual  removal  of  H-ions  leaves  in  the 
solution  an  excess  of  OH-ions  which  gives  a  basic  reaction 
toward  litmus.  (2)  Copper  sulphate  yields  the  ions  Cu+ 
and  SO4~~.  But  the  Cu-ions  combine  with  OH-ions  to 
form  copper  hydroxide  (Cu(OH)2),  which  dissociates  only 
to  a  very  slight  extent.  The  removal  of  OH-ions  leaves 
an  excess  of  H-ions  in  the  solution,  which  reacts  acid 
toward  litmus.  (3)  Acid  sodium  carbonate  (HNaCO3) 
yields  the  ions  Na+  andHCOs";  but  since  the  latter 
dissociates  to  only  a  very  slight  extent,  the  solution  is  not 
provided  with  an  excess  of  either  H-ions  or  OH-ions,  and 
therefore  has  a  neutral  reaction.  The  chemical  changes 
cited  in  this  paragraph  are  examples  of  a  general  phe- 
nomenon called  hydrolysis,  i.e.  a  chemical  change  (double 
decomposition)  due  to  the  interaction  of  certain  salts 
and  water.  The  behavior  of  the  final  solution  toward 
litmus  depends  upon  the  composition  of  the  salt,  as  the 
following  equations  show.  The  ordinary  chemical  equa- 
tion for  the  hydrolysis  of  sodium  carbonate  is  - 
Na2CO3  +  H20  =  NaHCO3  +  NaOH 

Sodium  Carbonate        Water        Acid  Sodium        Sodium 

Carbonate        Hydroxide 

The  ionic  equation  is  — 

2Na+  +  C03"  "  +  H+  +  OH"  -»  2Na+  +  HCO3~  +  OH" 


148 


CHEMISTRY 


The  corresponding  equations  for  the  hydrolysis  of  copper 
sulphate  are  — 


CuSO4 

Copper  Sulphate 


2H20  =       Cu(OH)2 

Water         Copper  Hydroxide 


H2SO4 

Sulphuric  Acid 


(See  Part  II,  Exp.  69.) 

166.  Electrolysis.  —  Electrolysis  is  the  series  of  changes 
accompanying  the  passage  of  an  electric  current  through 
a  solution  of  an  electrolyte,  i.e.  an  acid,  base,  or  salt. 
The  apparatus  in  which  the  operation  is  conducted  is 
called  an  electrolytic  cell  (Fig.  34  —  see  also  Figs.  22 

and  32).  The  electrolytic  cell 
generates  no  electric  current;  it 
receives  the  current  from  a  dy- 
namo or  a  battery. 

Two  illustrations  of  electrol- 
ysis  have   already   been   given 
-  (i)  a    dilute   sulphuric    acid 

solution.  (54),  which  yields  hy- 
Fig.  34.  -Electrolytic  Cell.    A    d  fa  h  d          d 

and  C  are  the  Electrodes,  E 

is  the  Solution  of  the  Elec-  gen  at  the  anode,  and  (2)  sodium 

chloride  solution  (75),  which 
yields  sodium  hydroxide  and 
hydrogen  at  the  cathode  and  chlorine  at  the  anode. 

Let  us  interpret  typical  illustrations  of  electrolysis  by 
the  theory  of  electrolytic  dissociation,  (a)  Hydrochloric 
acid  contains  hydrogen  ions  (H+)  and  chloride  ions  (Cl~). 
These  ions  are  formed  in  the  solution  as  soon  as  the  acid 
dissolves  and  before  the  electric  current  is  connected  with 
the  cell.  When  the  current  is  turned  on,  the  electrodes 
become  charged  with  electricity  —  the  anode  positively 
and  the  cathode  negatively.  Now  according  to  a  principle 


trolyte,  B  or  D  is  the  Bat- 
tery or  Dynamo. 


ACIDS,  BASES,  AND  SALTS 


149 


established  many  years  ago,  bodies  charged  with  like  kinds 
of  electricity  repel  each  other  and  bodies  charged  with 
unlike  kinds  attract  each  other.  Consequently  the  anions 
or  electro-negative  ions  move  toward  the  anode  or  elec- 
tro-positive electrode,  while  the  cations  or  electro-positive 
ions  move  toward  the  cathode  or  electro-negative  elec- 
trode, or  briefly,  "anions  to  anode,  cations  to  cathode." 
This  migration  of  the  ions,  as  it  is  called,  toward  their  re- 
spective electrodes  is  shown  diagrammatically  in  Fig.  35. 
As  soon  as  the  ions  reach 
their  electrodes  they  act  in 
accordance  with  another 
long-established  principle; 
that  is,  they  give  up  their 
electric  charges.  In  other 
words,  when  the  electro- 
positive cations  of  hydrogen 
touch  the  electro-negative 
cathode,  the  electric  charges 
of  the  ions  are  neutralized.  Electric  charges  lost  by  the 
cathode  are  constantly  renewed  by  the  battery  or  dy- 
namo; but  the  hydrogen  ions  once  deprived  of  their  elec- 
tric charges  do  not  regain  them  and  immediately  become 
ordinary,  uncharged  hydrogen  atoms,  which  combine  and 
escape  as  molecules  of  hydrogen  gas.  Similarly,  the  elec- 
tro-negative anions  of  chlorine  migrate  to  the  electro-posi- 
tive anode,  lose  their  charges,  become  chlorine  atoms,  and 
escape  as  chlorine  gas.  (b)  In  a  copper  sulphate  solution 
the  ions  are  copper  ions  (Cu++)  and  sulphate  ions(S04~~). 
When  this  solution  is  electrolyzed,  the  copper  ions  (Cu++) 
migrate  to  the  cathode,  lose  their  electric  charges,  become 
copper  atoms  (Cu),  and  adhere  as  metallic  copper  to  the 
cathode.  The  sulphate  ions  (SO4~~)  lose  their  electric 


.  35.  —  Migration  of  Ions  in  an 
Electrolytic  Cell. 


1 50  CHEMISTRY 

charges  at  the  anode  and  immediately  interact  with  the 
water  around  the  anode,  forming  sulphuric  acid  (H2SO4) 
and  oxygen  (O).  The  oxygen  escapes,  but  the  sulphuric 
acid  mingles  with  the  solution  and  dissociates  into  its 
ions,  (c)  Similarly,  in  a  solution  of  sodium  sulphate 
the  SO4-ions  undergo  the  change  described  in  (b),  while 
the  sodium  atoms  interact  with  the  water  around  the 
cathode  and  form  sodium  hydroxide  and  hydrogen.  (See 
Part  II,  Exps.  70,  71,  74.) 

The  so-called  electrolysis  of  water  is  interpreted  as  follows: 
Water  itself  does  not  conduct  electricity  to  an  extent  which  is 
comparable  with  the  behavior  of  an  electrolytic  solution,  because 
water  dissociates  only  inappreciably  and  gives  therefore  an  ex- 
ceedingly small  number  of  ions.  A  solution  of  sulphuric  acid  con- 
tains hydrogen  ions  (aH"1")  and  sulphate  ions  (SO4~~).  When  a 
^current  is  passed  through  this  solution,  hydrogen  ions  migrate  to 
the  cathode,  lose  their  electric  charges,  become  hydrogen  atoms, 
and  eventually  escape  as  hydrogen  gas;  the  SO4-ions  migrate  to 
the  anode,  lose  their  electric  charges,  and  interact  with  the  water 
to  form  sulphuric  acid  and  oxygen.  The  oxygen  escapes  as  a  gas, 
while  the  sulphuric  acid  dissociates  into  its  ions. 

We  may  now  re-define  or  describe  electrolysis  as  the 
migration  of  ions  induced  by  an  electric  current,  the  ions 
moving  to  their  respective  electrodes,  where  they  are 
transformed  into  atoms  or  atomic  groups,  which  escape 
wholly  or  in  part  as  elements  or  form  various  products  by 
interaction  with  the  water  of  the  solution. 

167.  Practical  Applications  of  Electrolysis.  —  Electrotyping  and 
electroplating  consist  in  depositing  a  thin  film  of  metal  upon  a 
surface.  The  process  of  electrotyping  is  substantially  as  follows: 
The  page  of  type,  for  example,  is  first  reproduced  in  wax.  This 
exact  impression  is  next  covered  with  powdered  graphite  to  make 
it  conduct  electricity.  The  coated  mold  is  then  suspended  as  the 
cathode  in  an  acid  solution  of  copper  sulphate;  the  anode  is  a 
plate  or  bar  of  copper.  When  the  current  is  passed,  electrolysis 


ACIDS,  BASES,  AND  SALTS 


occurs,  and  copper  is  deposited  upon  the  mold  in  a  film  of  any 
desired  thickness.  The  exact  copper  copy  is  stripped  from  the 
mold,  backed  with  metal,  and  used  instead  of  the  type  for 
printing.  The  process  of  electroplating  differs  from  electrotyping 
in  only  one  essential,  viz.,  in  electroplating  the  deposited  film  is 
not  removed  from  the  object. 

168.  Dissociation  of  Acids,  Bases,  and  Salts.  —  Dis- 
sociation, or  ionization  as  it  is  sometimes  called,  varies 
(i)  with  the  concentration  of  the  solution  —  being  less  in 
a  concentrated  than  in  a  dilute  solution,  and  (2)  with  the 
substance  itself,  (i)  Thus,  in  sodium  hydroxide  solution 
having  about  40  gm.  in  a  liter  the  per  cent  of  ionized  base 
is  73,  whereas  it  is  86  in  a  solution  having  4  gm.  in  a  liter. 
In  contrast  to  this  degree  of  dissociation,  acetic  acid 
solutions  having  respectively  60  gm.  and  6  gm.  in  a  liter 
contain  respectively  only  .4  and  1.3  per  cent  of  ionized 
acid.  The  approximate  per  cent  of  dissociation  of  cer- 
tain acids,  bases,  and  salts  in  solutions  of  the  same  rela- 
tive concentration  and  at  the  same  temperature  may  be 
seen  from  the  Table  of  Ionization. 

TABLE  or  IONIZATION 


Substance 

Ions 

Per  Cent  of 
Ionization 

Hydrochloric  Acid  ...    .                

H+,  Cl" 

GO 

Nitric  Acid                                                 .... 

H+,  NO3~ 

QO 

Sulphuric  Acid 

H+   HSO4~ 

60 

Acetic  Acid 

H+   H3C2O2~ 

I  "? 

Potassium  Chloride                 

K+,  Cl~ 

86 

Silver  Nitrate                                     

Ag+,  NO3~ 

86 

Barium  Chloride                                      

Ba++    2C1~ 

72 

Potassium  Hydroxide 

K+,  OH 

86 

Sodium  Hydroxide  

Na+,  OH~ 

86 

Ammonium  Hydroxide   

NH4+,  OH- 

1.4 

152 


CHEMISTRY 


(2)  The  simple  ions  formed  by  different  electrolytes 
are,  as  a  rule,  the  charged  atoms  or  groups  correspond- 
ing to  the  two  parts  of  the  compounds.  Thus,  we 
speak  of  nitric  acid  as  consisting  of  the  two  parts  H 
and  NOs;  similarly,  the  parts  of  sodium  nitrate  are  Na 
and  NO3,  and  of  sodium  hydroxide  are  Na  and  OH.  The 
corresponding  ions  are  H+,  NO3~,  Na+,  NO3~,  Na+, 
OH~.  The  ions  normally  formed  by  the  ionization  of 
acids,  bases,  and  salts  may  be  seen  in  the  Table  of  Ions. 
(See  Part  II,  Exps.  76,  77.) 

TABLE  OF  IONS 


Element  or 
Group 

Ion 

Element  or 
Group 

Ion 

Element  or 
Group 

Ion 

Hydrogen 

H+ 

Calcium     .  . 

Ca+  + 

Aluminium  . 

A1+  + 

Sodium  

Na+ 

Barium  

Ba+  + 

Antimony.  . 

Sb+  +  + 

Potassium 

K+ 

Copper  . 

Cu+  + 

Bismuth  .  .  . 

Bi+  +  + 

Silver  

A?+ 

Zinc  

Zn+  + 

Iron  (ic)  .  .  . 

Fe+  +  + 

Ammonium.  .  .  . 
Mercury  (ous)  . 

NH4+ 
Hg+ 

Magnesium  . 
Lead  

Mg+  + 
Pb+  + 

Tin  (ic)  .... 

Sn+  +  +  + 

Chlorine  
Sromine  
Iodine 

ci- 

Br- 
I- 

Iron  (ous)  .  .  . 
Mercury  (ic)  . 
Tin  (ous)  .  .  . 

Fe+  + 
Hg+  + 
Sn+  + 

Nitrate  

NO3~ 

Sulphate.  .  .  . 

so4-- 

Chlorate 

C1O3~ 

Sulphide  .... 

s— 

Hydroxyl  

OH- 

Carbonate  .  . 
Chromate.  .  . 
Bichromate  . 

co3-~ 

CrO4~~ 
Cr2O7~~ 

169.  Conclusion.  —  The  topics  discussed  in  the  pages 
immediately  preceding  show  that  the  theory  of  elec- 
trolytic dissociation  is  an  acceptable  explanation  of  many 
properties  of  dilute  solutions,  more  particularly  the  solu- 
tions of  acids,  bases,  and  salts. 


ACIDS,   BASES,   AND   SALTS  153 

EXERCISES 

1.  State  some  general  properties  of  solutions  of  acids,   bases,  and  salts 
which  differ  from  the  corresponding  properties  of  solutions  of  other  kinds 
of  substances. 

2.  State  the  theory  of  electrolytic  dissociation.    Illustrate  it. 

3.  Define  and  illustrate  the  terms  ion,  anion,  cation,  electrode,  anode, 
cathode,  positive  electrode,  negative  electrode. 

4.  Distinguish  between  an  atom  and  an  ion  of  potassium.    How  is  each 
represented? 

5.  Name  the  ions  in  a  solution  of  (a)  hydrochloric  acid,  (b)  sodium 
chlorate,    (c)  calcium   hydroxide,    (d)  sodium   nitrate,    (e)  zinc   sulphate. 

6.  In  what  respect  are  acids  alike?     Different?      Answer  for  bases. 

7.  Interpret  neutralization  by  the  theory  of  electrolytic  dissociation. 

8.  Discuss  salts  in  terms  of  the  theory  of  electrolytic  dissociation. 

9.  Define  electrolysis.    Interpret  two  or  more  cases  of  electrolysis  by 
the  theory  of  electrolytic  dissociation. 

10.  Write  ionic  equations  for  (a)  potassium  chloride  and  silver  sulphate 
form  silver  chloride  and  potassium  sulphate,  and   (b]   barium  nitrate  and 
sodium  sulphate  form  barium  sulphate  and  sodium  nitrate. 

11.  Write  the  name  and  formula  of  the  salt  formed  by  the  interaction 
of  calcium  hydroxide  and  (a)  nitrous  acid,  (6)  sulphurous  acid,  and  (c) 
chloric  acid. 

12.  Write  the  name  and  formula  of  the  sodium  salt  corresponding  to 
these  acids:    Carbonic,  chromic,  persulphuric,  manganic,  hypophosphorous, 
and  manganous.    Write  the  name  and  formula  of  the  corresponding  salts 
of    ammonium,    mercury  (ic),  lead,  magnesium,    aluminium,    silver,  and 
iron  (ous). 

PROBLEMS 

1.  Calculate  the  number  of  grams  of  hydrogen  in  40  gm.  of  (a)  sulphuric 
acid,  (b)  nitric  acid,  (c)  hydrochloric  acid. 

2.  Calculate  the  number  of  grams  of  hydroxyl  corresponding  to  (a)  75 
gm.  of  the  base  formed  by  Na,  (b)  35  gm.  of  A1(OH)3,  (c)  80  gm.  of  Ba(OH)2. 

3.  Calculate  the  weight  of  nitric  acid  needed  to  neutralize  27  gm.  of 
the  base  corresponding  to  these  metals:  sodium,  K,  calcium,  Ba,  NH,i. 

4.  Suppose  37.5  cc.  of  a  hydrochloric  acid  solution  neutralize  30  cc.  of 
a  sodium  hydroxide  solution,  and  that  each  cc.  of  the  sodium  hydroxide 
solution  contains  .003  gm.  of  the  base.    What  weight  of  hydrochloric  acid 
is  contained  in  the  acid  solution? 

5.  Two  solutions  are  thoroughly  mixed.    One  contains  75  gm.  of  sulphuric 
acid,  and  the  other  75  gm.  of  sodium  hydroxide.    What  will  be  the  weight 
of  the  compounds  (other  than  water)  in  the  final  solution? 


CHAPTER  XV 

CARBON  —  OXIDES  AND  CARBONATES  —  HYDROCARBONS  — 
CARBIDES  —  CYANOGEN 

170.  Occurrence    of    Carbon.  —  Uncombined    carbon 
occurs  as  diamond  and  graphite,  and  in  a  more  or  less 
impure  state  as  coal  and  similar  substances,  which  are 
included  in   the  term  amorphous  carbon.     Carbon  forms 
a  very  large  number  of  compounds,  natural  and  manu- 
factured.    Its  compounds  with  hydrogen  and  oxygen,  and 
with  nitrogen  also,  are  essential  parts  of  plants  and  ani- 
mals.     Meat,   starch,   fat,   sugar,   wood,   cotton,   paper, 
soap,  wool,  wax,  flour,   albumin,   rubber,  and  bone  are 
examples  of  natural  substances  which  contain  carbon.    It 
is  a  constituent  of  carbon  dioxide  and  of  carbonates,  such 
as  limestone,  chalk,  marble,  and  dolomite.     Illuminating 
gases,  gasoline,  and  kerosene  (and  other  substances  ob- 
tained from  petroleum),  and  turpentine  are  compounds  of 
carbon.    Among  the  manufactured  compounds  of  carbon 
are  dyestuffs,  many  medicines  and  perfumes,  soap,  and 
alcohol.     (See  Part  II,  Exp.  80.) 

171.  Diamond  is  pure   crystalline   carbon.     As  found 
in  nature,   diamonds    are    rough-looking    stones;    a  few 
are  crystals,  some  are  rounded  like  peas,  and  many  are 
irregular;   all    must  be  ground  into   special  shapes  and 
polished  to  bring  out  the  luster  and  make  them  sparkle 
(Fig.  36).     The  most  expensive  diamonds  are  colorless, 
transparent,  and  flawless.     Many  attractive  stones,  how- 
ever, are  slightly  tinted,  often  yellow,  rarely  pink  or  blue. 


CARBON 


155 


Diamond  has  the  high  specific  gravity  of  3.5,  and  is 
one  of  the  hardest  substances.  It  resists  the  action  of 
most  chemicals,  though  it  combines  with  oxygen  when  the 
two  elements  are  heated  together  to  a  high  temperature. 


Crystal  Rough 

Fig.  36.  —  Diamonds. 


Cut 


Diamonds  have  always  been  prized  as  gems  on  account  of  their 
beauty  and  permanency.  Besides  being  worn  as  jewels,  they  are 
used  to  cut  glass,  and  the  powder  and  splinters  (known  as  bort) 
are  used  to  grind  and  polish  diamonds  and  other  hard  gems. 
The  black,  impure  variety,  called  carbonado,  is  set  into  the  end 
of  the  "  diamond  drill,"  which  is  used  extensively  for  boring 
artesian  wells  and  drilling  hard  rocks. 

Diamond  is  carbon,  for  when  pure  diamond  is  burned  in 
oxygen,  the  only  product  is  carbon  dioxide.  Diamonds  have 
been  made  from  carbon.  The  French  chemist  Moissan  in  1893 
dissolved  pure  charcoal  in  melted  iron,  and  suddenly  cooled  the 
solution  in  water.  The  pressure  caused  the  cooling  carbon  to 
crystallize  into  diamond. 

The  largest  diamond  thus  far  known  was  called  the  Cullinan. 
It  was  found  in  South  Africa  in  1905,  and  weighed  3024!  carats 
(English,  or  1.37  Ib.  avoir.).1  It  has  been  cut  into  several  stones, 
which  are  among  the  crown  jewels  of  England. 

172.  Graphite  is  a  soft,  dark  lead-colored,  shiny  solid, 
which  is  smooth  and  greasy  to  the  touch.  Pure  graphite 
is  carbon;  it  is  sometimes  called  ''black  lead"  or  plum- 

1  The  carat  formerly  had  a  variable  weight.  In  the  United  States 
it  was  3^  Troy  grains  (or  205  milligrams).  The  metric  or  international 
carat,  which  is  200  milligrams,  was  adopted  July  i,  1913. 


156  CHEMISTRY 

bago,  because  it  was  formerly  supposed  to  contain  lead. 
Unlike  diamond,  graphite  is  a  good  conductor  of  electricity 
and  is  often  used  to  coat  molds  in  electrotyping.  It  is 
so  soft  that  it  readily  wears  away;  hence,  it  blackens  the 
fingers  and  leaves  a  black  mark  on  paper  when  drawn 
across  it.  This  property  is  indicated  by  the  name  graphite, 
which  is  derived  from  a  Greek  word  meaning  to  write. 
It  resembles  diamond  in  its  insolubility  in  liquids  at  the 
ordinary  temperature.  Its  specific  gravity  is  2.2,  being 
considerably  lighter  than  diamond.  It  changes  into 
carbon  dioxide  when  heated  intensely  in  oxygen;  but  it 
can  be  heated  to  a  very  high  temperature  in  the  air 
without  melting  or  undergoing  appreciable  change.  It  is 
used  to  make  stove  polish  and  protective  paints,  as  a 
lubricant  where  oil  might  be  decomposed  by  heat  or 
might  clog  machinery,  as  the  principal  ingredient  of 
graphite  crucibles  in  which  metals  are  often  melted,  and 
in  making  electrodes  for  electric  furnaces  and  electrolytic 
cells.  (See  Part  II,  Exp.  88.) 

Large  quantities  of  graphite  are  consumed  in  the  manufacture 
of  lead  pencils.  The  graphite  is  washed  free  from  impurities, 
ground  to  a  fine  powder,  mixed  with  more  or  less  clay,  and  then 
pressed  through  perforated  plates,  from  which  the  "  lead  "  issues 
in  tiny  rods.  These  are  dried,  cut  into  the  proper  lengths,  baked 
to  remove  all  traces  of  moisture,  and  then  inserted  in  the  wooden 
case.  Soft  pencils  contain  a  larger  proportion  of  graphite  than 
hard  ones. 

Graphite  occurs  abundantly,  the  famous  localities  being  Ceylon, 
Siberia,  and  Mexico.  Considerable  graphite  is  manufactured  at 
Niagara  Falls  by  heating  a  special  grade  of  coal  or  of  coke  out  of 
contact  with  air  in  an  electric  furnace.  The  process  is  electro- 
thermal, i.e.  the  chemical  change  is  brought  about  by  the  heat 
produced  by  passing  an  electric  current  through  the  materials 
(compare  200).  Articles  of  almost  any  size  can  now  be  made  of 
graphite,  which  is  so  compact  that  it  can  be  further  shaped  by 


CARBON  157 

tools.     This  kind  of  graphite  is  especially  suitable  for  the  elec- 
trodes used  in  electrolytic  and  electrothermal  apparatus. 

173.  Amorphous  Carbon  includes  all  varieties  of  coal 
and  charcoal,  lampblack,  coke,  and  gas  carbon.     They 
are  the  non-crystalline  forms  of  impure  carbon.    The  word 
amorphous  means    literally    "  without    form,"   and  it  is 
often  used  to  designate  soft,  powdery,  and  uncrystallized 
substances. 

174.  Coal. --There  are  two  principal  kinds  of  coal, 
though  several  varieties  of  each  are  common  articles  of 
commerce,     (i)  Bituminous  or  soft  coal  contains  about 
70  per  cent  of  carbon.    It  burns  with  a  smoky  flame,  and 
is  used  to  make  illuminating  gas,  coke,  and  as  a  fuel  for 
steam.     (2)  Anthracite  or  hard  coal  contains  90  per  cent 
or  more  of  carbon.     It  ignites  with  difficulty,  burns  with 
little  or  no  flame,  and  produces  considerable  heat.    It  is 
used  mainly  for  domestic  purposes, — heating  and  cooking, 

-  especially  in  the  eastern  United  States.  Some  varieties 
of  coal,  such  as  lignite  or  brown  coal,  contain  only  a  small 
proportion  of  carbon,  as  low  as  20  per  cent.  Yet  they 
are  used  as  fuel  in  some  localities. 

Besides  carbon,  coal  contains  moisture,  ash  or  mineral 
matter,  and  volatile  matter  which  consists  mainly  of 
compounds  of  hydrogen  and  carbon.  Coal  is  used  as  fuel 
because  when  the  carbon  and  gases  are  burned,  heat  is 
readily  liberated.  The  heat  producing  value  of  coal  is 
measured  in  calories  per  gram,  i.e.  the  number  of  calories 
liberated  when  i  gram  is  burned  freely  (compare  164). 
Thus,  a  well  known  commercial  variety  of  anthracite  coal 
liberates  about  7724  calories  and  of  bituminous  coal  8768 
calories.  Coal  is  often  purchased  on  the  basis  of  its 
calorific  value  (182).  (See  Part  II,  Exp.  81.) 


158 


CHEMISTRY 


Ages  ago  the  vegetation  was  exceedingly  dense  and  luxuriant 
upon  land  slightly  raised  above  the  sea.  In  process  of  time  this 
vegetation  decayed,  accumulated,  and  slowly  became  covered  with 
sand,  mud,  and  water.  This  vegetable  matter  was  then  slowly 


Fig.  37.  —  Section  of  Part  of  the  Earth's  Crust  near 
Mauch  Chunk,  Perm.,  Showing  Layers  of  Coal. 

changed  into  more  or  less  impure  solid  carbon,  moisture,  and 
gaseous  and  liquid  compounds  called  hydrocarbons.  The  geologi- 
cal and  chemical  changes  were  repeated,  and  as  a  result  we  find 
in  the  earth  layers  or  seams  of  carbonaceous  matter  varying  in 
thickness  and  composition  (Fig.  37).  These  are  the  coal  beds. 
Coal  beds  contain  proofs  of  their  vegetable  origin,  viz.  impressions 


Fig.  38.  —  Fossil  Found  in 
a  Coal  Bed. 


Fig.  39.  —  Section  of  Coal  as  Seen 
through  a  Microscope. 


of  vines,  stems,  and  leaves  of  plants,  and  similar  vegetable  sub- 
stances (Fig.  38).  A  thin  section  of  coal  examined  through  a 
microscope  reveals  a  distinct  vegetable  structure  (Fig.  39). 

175.    Charcoal  is  obtained   by  heating  wood,   bones, 
and  other  organic  matter  in  closed  vessels,  or  by  partially 


CARBON,  159 

burning  them  in  the  air.  The  process  consists  essentially 
in  driving  off  the  volatile  matter.  Besides  carbon,  char- 
coal contains  mineral  matter.  (See  Part  II,  Exp.  82.) 

176.  Wood  Charcoal  is  a  black,  brittle  solid,  and  often 
retains  the  form  of  the  wood  from  which  it  is  made.  It 
is  insoluble,  though  its  mineral  impurities  may  be  removed 
by  acids.  It  burns  without  flame  or  much  smoke,  and 
leaves  a  white  ash  which  consists  of  mineral  substances. 
The  compact  varieties  conduct  heat  and  electricity,  but 
porous  charcoal  is  a  poor  conductor.  It  resists  the  action 
of  moisture  and  many  chemicals;  hence  fence  posts, 
telegraph  poles,  and  wooden  piles  are  often  charred  before 
being  put  into  the  ground.  Most  varieties  are  very  porous, 
and  are  excellent  absorbers  of  gases.  Sewers,  cisterns, 
and  foul  places  are  sometimes  purified  by  charcoal.  It 
will  also  absorb  colored  substances  from  solutions.  This 
is  especially  true  of  animal  charcoal  (177).  Foul  air 
and  water  may 
be  partially  pu- 
rified by  char- 
coal, which 
forms  the  es- 
sential part  of 
many  water  fil- 
ters in  houses. 
Charcoal  used 

r  ,  Fig.  40.  —  Wood  Arranged  for  Burning  into  Charcoal. 

lor  sucn  a  pur- 

pose,  however,  must  be  frequently  renewed  or  often  heated 
to  redness;  otherwise  it  becomes  clogged  and  contami- 
nated. Large  quantities  of  charcoal  are  used  as  a  fuel 
and  in  the  manufacture  of  steel  and  of  gunpowder. 

Wood  charcoal  is  made  either  in  charcoal  pits  or  in  large  iron  fur- 
naces.    Where  wood  is  plentiful,  it  is  loosely  piled  into  the  shape 


160  CHEMISTRY 

shown  in  Fig.  40,  and  covered  with  turf.  The  wood  is  lighted, 
and  as  it  slowly  burns  care  is  taken  to  regulate  the  supply  of 
air  so  that  the  wood  will  smolder  but  not  be  entirely  consumed. 
The  volatile  matter  escapes  and  charcoal  remains.  Much  charcoal 
is  made  by  heating  wood  in  closed  furnaces,  no  air  whatever  being 
admitted.  By  this  method,  which  is  called  dry  or  destructive 
distillation  because  the  wood  is  destroyed,  the  yield  of  charcoal  is 
about  30  per  cent  and  all  the  volatile  matter  is  saved.  More  or 
less  charcoal  is  obtained  by  heating  any  compound  of  carbon,  e.g. 
sugar  or  starch,  the  charring  being  a  test  for  carbon. 

177.  Animal  Charcoal  or  Bone  Black  is  made  by  heat- 
ing bones  and  animal  refuse  in  a  closed  vessel.     The 
animal  charcoal  from  bones  contains  only  about  10  per 
cent  of  carbon,  but  this  carbon  is  distributed  throughout 
the  porus  mineral  matter  of  the  bone,  which  is  largely 
calcium  phosphate.    Under  the  name  of  ivory  black,  ani- 
mal charcoal  is  used  as  a  pigment,  especially  in  making 
shoe-blacking.      It  is   extensively   used   to   remove   the 
color  from  sugar  sirups,  oils,  and  other  liquids  colored  by 
organic  matter. 

178.  Coke  is  made  by  expelling  the  volatile  matter 
from  bituminous  coal,  somewhat  as  charcoal  is  made  from 
wood.     It  is  left  in  the  retorts  when  coal  is  distilled  in 
the  manufacture  of  coal  gas  (205).    On  a  large  scale  it  is 
made  by  heating  a  special  grade  of  soft  coal  in  huge  brick 
ovens,  shaped  like  a  beehive,  from  which  air  is  excluded 
after  combustion  begins.    Sometimes  the  coke  is  made  in 
closed  furnaces  constructed  so  as  to  save  the  by-products, 

—  ammonia,  tar,  organic  compounds,  and  combustible 
gases.  Coke  is  a  grayish,  porous  solid,  harder  and 
heavier  than  charcoal.  It  burns  with  no  smoke  and  a 
feeble  flame. 

179.  Gas   Carbon  is  a  variety  of  amorphous  carbon  which  is 
deposited  on  the  inside  of  the  retorts  used  in  the  manufacture  of 


CARBON  161 

illuminating  gas  (205).  It  is  a  black,  heavy,  hard  solid,  and  is 
almost  pure  carbon.  It  is  a  good  conductor  of  electricity,  and 
is  extensively  used  for  the  manufacture  of  the  carbon  rods  of 
electric  lights  and  for  plates  of  electric  batteries. 

180.  Lampblack  is  prepared  by  burning  oil  or  oily  substances  rich 
in  carbon  in  a  limited  supply  of  air.      The  dense  smoke,  which  is 
mainly  finely  divided  carbon,  is  passed  through  a  series  of  conden- 
sing chambers,  where  it  is  collected  upon  coarse  cloth  or  a  cold  sur- 
face.     Lampblack  is  one  of  the  purest  forms  of  amorphous  carbon, 
and  it  is  used  in  making  printer's  ink  and  certain  black  paints. 

181.  Allotropism. — Diamond,  graphite,  and  pure  amor- 
phous carbon,  though  having  essentially  different  prop- 
erties, are  different  modifications  of  the  element  carbon. 
They  can  be  changed  into  one  another.     Each  burns  in 
oxygen  and    the    product  is  carbon  dioxide.      Further- 
more, the  same  weight  of  each  forms  the  same  weight 
of  carbon  dioxide.     Elements  which  exist  in  two  or  more 
distinct  modifications  are  called  allotropic,  and  the  prop- 
erty of  assuming  more  than  one  elementary  modification 
is  called  allotropism  or  allotropy  (from  Greek  words  mean- 
ing another  form). 

182.  Chemical  Properties  of  Carbon.  —  Carbon  does  not  interact 
with  acids  or  bases,  but  it  is  readily  oxidized  by  salts  which  liber- 
ate oxygen,  such  as  potassium  nitrate  and  potassium  chlorate.      It 
unites  directly  with  certain  elements,  though  the  temperature  must 
usually  be  raised  to  bring  about  combination.      Thus,  with  oxygen 
it  forms  carbon  dioxide  and  carbon  monoxide,  and  with  sulphur, 
carbon  disulphide.      Carbon  unites  with  many  metals  and  some 
non-metals  at  the  high  temperatures  produced  in  the  electric  fur- 
nace, thereby  forming  carbides   (199-201).      Considerable  heat  is 
developed  during  the  uniting  of  carbon  with  oxygen;  the  use  of 
various  forms  of  carbon  and  carbonaceous  substances  as  fuel  is 
based  on  this  fact.      The  thermal  equation  for  the  transformation 
of  carbon  (in  the  form  of  charcoal)  into  carbon  dioxide  is: — 

C     +     O2    =        CO2      +    97,000  calories 

Carbon        Oxygen        Carbon  Dioxide 


1 62  CHEMISTRY 

Carbon  reduces  oxides,  a  simple  illustration  being  the   reduction 
of  copper  oxide,  thus:  — 

2CuO      +     C     =        CO2      +     2Cu 

Copper  Oxide        Carbon        Carbon  Dioxide        Copper 

Different  varieties  of  carbon  are  used  as  reducing  agents,  especially 
coke  in  the  manufacture  of  iron  from  ferric  oxide  (Fe2O3). 


Oxides  of  Carbon,  Carbonic  Acid,  Carbonates 

183.    Occurrence  and  Formation  of  Carbon  Dioxide.  - 

The  occurrence  of  carbon  dioxide  in  the  atmosphere  and 
in  many  natural  waters  has  already  been  mentioned  (128, 
43).  Carbon  dioxide  is  one  product  of  ordinary  com- 
bustion, respiration  of  animals;  fermentation,  and  decay. 
(See  Part  II,  Exp.  89.) 

Ordinary  combustion  is  a  chemical  combining  of  carbon 
and  oxygen.  The  equation  for  this  change  is:  - 

C     +     O2       =        CO2 

Carbon        Oxygen        Carbon  Dioxide 

Carbon  dioxide  is  formed  by  the  free  combustion  of  such 
common  substances  as  wood,  coal,  charcoal,  coke,  oils, 
waxes,  cotton,  bone,  starch,  sugar,  meat,  bread,  alcohol, 
camphor,  and  illuminating  gas.  This  fact  may  be  shown 
by  bubbling  the  smoke  from  these  burning  substances 
through  calcium  hydroxide  solution  (128). 

The  continuous  oxidation  of  the  tissues  of  the  body 
produces  carbon  dioxide  (186) .  And  if  we  exhale  the  breath 
through  a  glass  tube  into  calcium  hydroxide  solution,  the 
carbon  dioxide  which  is  in  the  breath  turns  the  solution 
milky  —  the  usual  test  for  carbon  dioxide.  The  equation 
for  the  change  is:  - 

CO2         +       Ca(OH)2  CaCO3     +      H2O 

Carbon  Dioxide        Calcium  Hydroxide        Calcium  Carbonate 


CARBON 


163 


Many  kinds  of  organic  matter  ferment,  especially  those 
containing  certain  varieties  of  sugar.  By  alcoholic  fer- 
mentation the  sugar  changes  into  carbon  dioxide  and 
alcohol  (235),  thus:- 

C6Hi2O6  2CO2     +     2C2H6O 

Sugar  Carbon  Dioxide  Alcohol 

Large  quantities  of  carbon  dioxide  are  liberated  in  the 
fermenting  vats  of  breweries. 

184.  The  Preparation  of  Carbon  Dioxide  is  usually 
accomplished  by  the  interaction  of  a  carbonate  and  an 
acid.  Calcium  carbonate  (limestone  or  marble)  and  hydro- 
chloric acid  are  ordinarily  used.  (See  Part  II,  Exp.  83.) 
The  equation  for  the  change  is  :  — 


CaC03  +      2HC1      =    C02 

Calcium        Hydrochloric        Carbon 
Carbonate  Acid  Dioxide 


CaCl2  +  H2O 

Calcium        Water 
Chloride 


185.  Properties  of  Carbon  Dioxide.  —  This  gas  has 
many  important  properties  besides  those 
already  mentioned  (128).  It  has  a  slight 
taste  and  odor,  but  no  color.  It  is  heavier 
than  air  (1.5  to  i),  and  a  liter  under  stand- 
ard conditions  weighs  1.977  gm.  At  ordi- 
nary temperatures  and  pressures,  water 
dissolves  its  own  volume  of  carbon  diox- 
ide. Under  increased  pressure  more  gas 
dissolves,  which  escapes  when  the  pressure 
is  removed.  Hence  soda  water,  which  is 
made  by  forcing  carbon  dioxide  into  water, 
effervesces  when  drawn  from  a  soda  foun- 
tain or  siphon  (Fig.  41).  Many  natural  waters  and  manu- 
factured beverages  effervesce  for  the  same  reason,  the 


Fig.  41. — Siphon. 


164 


CHEMISTRY 


gas  often  escaping  rapidly  in  bubbles.     This  gas  can  be 
readily  liquefied. 

Carbon  dioxide  does  not  burn,  but  extinguishes  burn- 
ing substances.  Instead  of  the  gas  itself,  a  saturated 
solution  is  frequently  used.  The  solu- 
tion is  prepared,  as  needed,  in  portable 
fire  extinguishers  and  chemical  engines 
by  the  interaction  of  sulphuric  acid  and 
sodium  bicarbonate  (HNaCOs).  The  or- 
dinary fire  extinguisher  contains  a  solu- 
tion of  sodium  bicarbonate  and  a  loosely 
stoppered  bottle  of  sulphuric  acid;  upon 
inverting  the  tank,  the  stopper  of  the 
acid  bottle  falls  out,  the  two  liquids  mix, 
and  the  pressure  of  the  generated  gas 
forces  the  saturated  solution  of  carbon 
Fig.  42.  —  Portable  dioxide  out  of  the  nozzle  of  the  extin- 

£5X5  guisher  (Fig- 42)- 

Showing  Stop-  186.  Relation  of  Carbon  Dioxide  to 
Life.  —  Carbon  dioxide  is  not  poisonous, 
though  the  presence  of  a  small  quantity 
in  the  air  is  objectionable.  As  already 
stated  the  carbon  dioxide  that  is  exhaled  from  our  lungs  is 
one'  of  the  products  formed  by  the  oxidation  of  the  tissues 
of  the  body,  new  tissue  itself  being  formed  from  the  food 
(18,  128,  see  also  Food  and  Nutrition,  Chapter  XVII). 
The  carbon  needed  for  the  rebuilding  of  tissue  is  supplied 
by  starch  and  other  substances.  Carbon  dioxide  is  a 
waste  product  of  animal  life.  On  the  other  hand,  carbon 
dioxide  is  an  essential  food  of  plants.  Through  their  leaves 
and  other  green  parts  they  absorb  carbon  dioxide  from 
the  atmosphere,  decompose  it,  reject  part  of  the  oxygen, 
and  store  up  the  carbon  in  the  form  of  complex  com- 


pered Acid  Bot- 
tle  in  Original 
Position. 


CARBON 


165 


pounds,  such  as  starch  ((C6Hi0O5)x).    The  sunlight  and  the 

green  coloring  matter  (called  chlorophyl)  aid  the  plant 

in  the   formation  of   these  compounds.     The  relation  of 

carbon  dioxide  to  plants  and  animals 

is  impressive.    Plants  absorb  carbon 

dioxide  and  transform  it  into  starch, 

while    animals    eat    starch    as    food, 

assimilate  it,  and  oxidize  the  carbon 

to  carbon  dioxide,  which  is   exhaled 

into  the   atmosphere    ready   for   the 

plants  again,  and  so  on. 

The  fact  that  plants  take  up  carbon  di- 
oxide and  reject  oxygen  can  be  readily  illus- 
trated, as  shown  in  Fig.  43.  Fresh  green 
leaves  are  put  into  the  large  vessel,  which  is 
then  completely  filled  with  water  saturated 
with  carbon  dioxide.  The  stopper  with  its 
funnel  is  pushed  in  to  exclude  the  air,  the 
funnel  is  partly  filled  with  the  same  liquid, 
and  the  test  tube  is  filled  and  arranged  as  FiS-  43-  —  Experiment 


Showing  the  Libera- 
tion of  Oxygen  by 
Plants. 


shown  in   the  figure.      On  exposure  to  the 

sunlight  for  several  hours,  a  gas  collects  in 

the  test  tube.      The  usual  test  shows  that  the 

gas  is  oxygen.      If  necessary,  a  second  experiment  may  be  done  to 

show  that  no  oxygen  is  produced  when  the  vessel  contains  water 

freed  from  carbon  dioxide  and  other  gases. 

The  significant  relation  of  carbon  dioxide  and  oxygen 
to  plants  and  animals,  which  is  often  spoken  of  as  the 
cycle  of  carbon  and  oxygen,  is  shown  diagramatically  in 
Fig.  44. 

187.  Carbonic  Acid.  —  Carbon  dioxide  gas  is  often  called  carbonic 
acid  gas,  or  simply  carbonic  acid;  the  latter  term  is  incorrect  when 
applied  to"  carbon  dioxide  and  should  be  used  only  as  the  name  of 
the  compound  H2CO3.  If  carbon  dioxide  is  passed  into  water,  it 


i66  CHEMISTRY 

combines  to  some  extent  with  the  water  and  forms  an  acid,  which 
is  carbonic  acid  (H2CO3).  Carbon  dioxide  is  sometimes  called 
carbonic  anhydride  to  emphasize  the  fact  that,  like  oxides  of  other 
non-metals,  it  unites  directly  with  water  and  thereby  forms  an 
acid.  A  solution  of  carbonic  acid  is  very  weak,  i.e.  it  ionizes  only 

Carbon  dioxide 
in  the  air 


Animals 
Animals^ 


Carbon  dioxide 
in  the  air 

Fig.  44. —  Cycle  of  Carbon  (A)  and  Oxygen  (B). 

slightly,  its  ions  being  H+  and  HCO3~  (1'68).  Carbonic  acid  has 
never  been  obtained  free,  and  is  so  unstable  that  it  easily  breaks 
up  by  gentle  heat  into  carbon  dioxide  and  water,  thus:  — 

H2CO3  =  CO2  +  H2O 

188.  Carbonates  are  salts  of  carbonic  acid.  They  are 
common  substances,  and,  unlike  the  acid,  are  stable 
under  ordinary  conditions.  Many  have  important  uses 
(see  Calcium  and  Sodium  Carbonates).  The  most  abun- 
dant natural  carbonates  are  those  of  calcium,  magnesium, 
and  iron.  Immense  quantities  of  sodium  and  potassium 
carbonates  are  manufactured. 

Many  carbonates  are  insoluble  in  water,  e.g.  calcium 
carbonate,  the  test  for  carbon  dioxide  depending  on  this 
fact.  Others,  e.g.  sodium  and  potassium  carbonates,  are 
very  soluble.  There  are  two  classes  of  carbonates,  the 
normal  and  the  acid.  Normal  sodium  carbonate  is 
Na2CO3,  and  acid  sodium  carbonate  is  HNaCOs.  The 
latter  is  often  called  sodium  bicarbonate.  Notmal  cal- 
cium carbonate  is  CaCO3,  and  acid  calcium  carbonate  is 


CARBON  167 

the  latter  is  soluble  in  water  and  is  easily 
decomposed  by  heat  into  normal  calcium  carbonate.  Acid 
calcium  carbonate  is  readily  formed  from  the  normal 
carbonate.  When  carbon  dioxide  is  passed  into  water 
containing  insoluble  normal  calcium  carbonate  in  sus- 
pension, soluble  acid  calcium  carbonate  is  formed, 
thus:- 

H2CO3      +         CaCO3  H2Ca(CO3)2 

Carbonic  Acid         Calcium  Carbonate        Acid  Calcium  Carbonate 

Now  when  this  solution  of  acid  calcium  carbonate  is 
heated,  the  decomposition  takes  place,  thus:  - 

H2Ca(CO3)2  CaCO3      +      CO2    +   H2O 

Acid  Calcium  Carbonate     Calcium  Carbonate        Carbon         Water 

Dioxide 

Since  many  underground  waters  contain  carbon  dioxide, 
these  waters  dissolve  the  limestone  (CaCO3)  over  which 
they  pass,  forming  "hard"  water.  When  the  dissolved 
acid  calcium  carbonate  is  decomposed  by  heat  or  in  some 
other  way,  the  calcium  carbonate  is  reprecipitated.  (See 
Stalactites  (under  Calcium  Carbonate),  Natural  Waters, 
and  Hardness  of  Water.) 

189.   Carbon    Monoxide    is    formed    when    carbon   is 
burned  in  a  limited  supply  of  air,  thus:  — 

2C         +         02  2CO 

Carbon  Oxygen  Carbon  Monoxide 

If  carbon  dioxide  is  passed  over  heated  charcoal,  the 
product  is  carbon  monoxide.  That  is,  carbon  reduces 
carbon  dioxide  to  carbon  monoxide,  the  equation  for  the 
change  being:  — 

CO2  +  C  =  2CO 


i68 


CHEMISTRY 


This  chemical  change  takes  place  in  every  coal  fire  (Fig. 
45).  The  oxygen  of  the  air  entering  at  the  bottom  of  the 
fire  unites  with  the  carbon  to  form  carbon  dioxide;  the 
latter  gas  in  passing  through  the  hot  carbon  of  the  fire  is 
reduced  to  carbon  monoxide.  Some  of  the  carbon  monox- 
ide escapes,  while  some  burns  with  a  flickering  bluish 
flame  on  the  top  of  the  fire  and  forms  carbon  dioxide.  If 
steam  is  passed  through  a  very  hot  hard  coal  or  coke 


2  CO  +  Oo  =  2  CO2 

C02  +  C  =2  CO 

2  C  +  2  O2  =  2  CO2 


JL 


\\ 


Fig.  45. —  Essential  Chemical  Changes  during  Combustion  in  a  Coal  Fire. 

fire,  a  mixture  of  carbon  monoxide  and  hydrogen  is 
formed;  this  mixture  enriched  by  vapor  from  oils  is 
known  as  water  gas  (206-7).  If  steam  and  air  are 
together  passed  through  the  hot  carbon,  the  gaseous 
product  contains  nitrogen  and  some  carbon  dioxide  be- 
sides carbon  monoxide  and  hydrogen;  it  is  called  fuel 
or  producer  gas,  and  since  it  is  easily  made  and  lib- 
erates considerable  heat  in  burning,  it  is  used  extensively 
as  a  fuel  in  industrial  processes,  e.g.  in  making  open 
hearth  steel. 

Carbon  monoxide  is  usually  prepared  in  the  laboratory 
by  heating  a  mixture  of  oxalic  acid  and  sulphuric  acid  in 
a  flask,  and  collecting  the  gaseous  product  over  water. 
The  oxalic  acid  decomposes  thus :  - 


CARBON  169 

C2H2O4  CO  +        CO2       +     H2O 

Oxalic  Acid        Carbon  Monoxide        Carbon  Dioxide 

The  carbon  dioxide  may  be  removed  by  passing  the 
mixed  gases  through  a  solution  of  sodium  hydroxide. 
(See  Part  II,  Exp.  91.) 

Carbon  monoxide  is  a  gas  without  color,  odor,  or  taste, 
and  is  only  slightly  soluble  in  water.  It  burns  with  a 
bluish  flame,  forming  carbon  dioxide,  thus :  — 

2  CO         +        O2  2CO2 

Carbon  Monoxide  Carbon  Dioxide 

Carbon  monoxide  is  extremely  poisonous,  and  it  is  very 
dangerous  because  the  lack  of  odor  prevents  its  detection 
in  time  to  escape  its  stupefying  effect.  Many  deaths 
have  been  caused  by  breathing  air  containing  it.  Carbon 
monoxide  impoverishes  the  blood  by  forming  a  compound 
with  one  of  its  constituents,  and  persons  who  have 
been  poisoned  by  this  gas  cannot  usually  be  revived  by 
air,  as  in  the  case  of  suffocation  by  carbon  dioxide.  It 
is  an  ingredient  of  ordinary  illuminating  gas,  and  care 
should  always  be  taken  to  prevent  the  escape  of  illuminat- 
ing gas  (as  well  as  the  gas  from  a  coal  stove  or  furnace) 
into  rooms  occupied  by  human  beings.  The  pulmotor  is 
often  used  to  revive  persons  who  have  been  overcome  by 
illuminating  gas  (19).  At  a  high  temperature  carbon 
monoxide  readily  reduces  oxides,  and  it  is,  therefore,  an 
important  agent  in  .the  reduction  of  iron  ores  in  the  blast 
furnace.  This  reaction  may  be  represented  thus:  - 
Fe2O3  +  3CO  =  2Fe  +  3CO2 

Iron  Oxide     Carbon  Monoxide        Iron        Carbon  Dioxide 

Carbon  monoxide,  which  is  sometimes  called  carbonic  oxide, 
forms  no  acid  and  no  salts.  It  does  not  turn  calcium  hydroxide 
milky,  thus  being  distinguished  from  carbon  dioxide.  Its  blue  flame 
serves  as  a  test  to  distinguish  it  from  other  combustible  gases. 


i  yo  CHEMISTRY 

Hydrocarbons 

190.  Hydrocarbons    are    compounds    of    carbon    and 
hydrogen.     There   are   about   two   hundred.     They   are 
found  in  petroleum  and  its  products  (gasoline,  kerosene, 
naphtha,  lubricating  oils,  paraffin  wax,  etc.)  ,  coal  tar,  coal 
gas,  and  natural  gas.     They  are  prepared  by  the  distilla- 
tion of  petroleum  and  coal  tar.     All  hydrocarbons  burn 
readily,  water  and  carbon  dioxide  being  the  products  of 
combustion.     Some  form  highly  explosive  mixtures  with 
air,  and  great  care  should  be  taken  in  using  naphtha  and 
gasoline,  which  contain  volatile  hydrocarbons. 

191.  Methane,   CH4,  is  found  in  coal  mines.     It  is 
called  fire   damp   by  miners.    It   is   formed   in   marshy 
places  by  the  decay  of  vegetable  matter  under  water,  and 
is  sometimes  called  marsh  gas.     The   illuminating   gas 
manufactured   from    coal    (208)    is   about   35    per   cent 
methane,  while  natural  gas,  which  issues  from  the  earth 
in  Pennsylvania,  West  Virginia,  Ohio,  and  other  localities 
contains  about  90  per  cent  of  methane. 

Methane  is  usually  prepared  in  the  laboratory  by  heating  a 
mixture  of  sodium  acetate,  sodium  hydroxide,  and  lime.  It  may 
also  be  prepared  by  the  interaction-  of  aluminium  carbide  and 
water,  thus:  — 


A14C3         +        i2H2O    =     3CH4        +       4A1(OH), 

Aluminium  Carbide  Water  Methane  Aluminium  Hydroxide 

Methane  is  a  gas  without  color,  taste,  or  odor.  It  burns 
with  a  pale,  luminous,  rather  hot  flame;  it  is  the  chief  heat 
producing  substance  in  natural  gas,  which  is  extensively 
used  in  natural  gas  regions  for  heating  houses  and  as  fuel 
in  many  industries,  e.g.  making  steel,  glass,  brick,  and 
pottery.  A  mixture  of  methane  with  oxygen  or  air  ex- 
plodes violently  when  ignited  by  a  spark  or  flame.  Ter- 


CARBON 


171 


rible  disasters  sometimes  occur  in  bituminous  coal  mines 
from  this  cause.  When  methane  burns  or  explodes, 
carbon  dioxide  and  water  are  formed,  thus:  - 

CH4        -h        2O2  CO2        +        2H2O 

Methane  Oxygen  Carbon  Dioxide  Water 

The  carbon  dioxide,  called  choke  damp  or  black  damp  by 

the  miners,  often  suffocates   those  who 

escape  from  the   explosion.      (See   Oxy- 
gen Helmet,  19.) 

The  miner's  safety  lamp  was  invented 

in  1815  by  Davy  to  prevent  explosions. 

It  is  essentially  an  oil  lamp  surrounded 

by  a  cylinder  of  fine  wire  gauze  instead  of 

the  usual  chimney  (Fig.  46) .    When  taken 

into  a  mine  where  there  are  explosive 

gases,  e.g.  methane,  the  gas  enters  the 

lamp  and  burns  inside,  but  the   flame 

within  does  not  ignite  the  gases  outside 

because  the  wire  gauze  keeps  them  cooled 

below  their  kindling  temperature.    When 

miners  notice  changes  in  the  lamp  flame, 

they  usually  seek  a  safe  place.    Modified  forms  of  the 

Davy  lamp  are  used  at 
the  present  time. 

The  principle  on  which  the 
safety  lamp  depends  can  be 
shown  by  a  simple  experiment. 
If  a  piece  of  wire  gauze  is  held 
a  few  inches  above  a  Bunsen 
burner  and  the  gas  is  turned 
on  and  lighted  above  the 
gauze,  the  flame  will  not  pass 
through  the  gauze  unless  the 
latter  is  hot  (Fig.  47).  (See  Part  II,  Exp.  92.) 


Fig.  46.  —  A  Form 
of  Davy's  Safety 
Lamp  (gauze  cut 
away  to  show 
inside). 


Fig.  47.  —  Experiment  Illustrating  the 
Principle  of  Davy's  Safety  Lamp. 


172  CHEMISTRY 

Two  important  substitution  products  of  methane  are  chloroform 
(CHC13)  and  iodoform  (CHI3).  Chloroform  is  a  heavy,  sweet, 
volatile  liquid,  and  iodoform  is  a  yellow,  disagreeable  smelling  solid. 
Chloroform  is  used  as  an  anaesthetic,  i.e.  to  produce  unconsciousness 
in  surgical  operations.  Iodoform  is  used  as  an  antiseptic  dressing 
for  cuts  and  wounds. 

192.  Ethylene,  C2H4,  is  usually  prepared  in  the  labo- 
ratory by  heating  a  mixture  of   concentrated  sulphuric 
acid  and  ethyl  alcohol.      Ethylene  is  a  colorless  gas  with 
a  pleasant  odor.      It  burns  with  a  yellow  flame,  and  is 
one  of  the  illuminating  constituents  of  coal  gas   (208). 
The    complete    combustion   of    ethylene    is    represented 
thus:- 

C2H4      +      302      =      2C02      +      2H20 

Ethylene  Carbon  Dioxide          Water 

193.  Acetylene,  C2H2,  is  prepared  commercially  and  in 
the  laboratory  by  the  interaction  of  calcium  carbide  and 
water,  thus :  — • 

CaC2     +     2H20     =     C2H2     +     Ca(OH)2 

Calcium  Carbide         Water  Acetylene      Calcium  Hydroxide 

Acetylene  is  a  colorless  gas,  and  if  prepared  from  cal- 
cium carbide  the  odor  is  unpleasant.  It  is  poisonous  if 
breathed  in  large  quantities,  but  much  less  dangerous 
than  gases  containing  carbon  monoxide.  Acetylene  burns 
in  the  air  with  a  luminous,  smoky  flame.  But  when  con- 
siderable air  is  mixed  with  the  gas  as  the  latter  issues  from 
a  small  opening,  the  mixture  burns  with  a  brilliant,  white 
flame,  which  does  not  smoke.  It  is  used  as  an  illuminant 
in  localities  where  ordinary  illuminating  gas  is  not  avail- 
able; it  is  also  used  in  automobile  lights.  With  a  proper 
burner  the  combustion  of  acetylene  is  complete,  and  may 
be  represented  thus:  - 


CARBON 


2C2H2 

Acetylene 


=     4C02     +     2H2O 

Oxygen      Carbon  Dioxide         Water 


Ordinary  gas  burners  cannot  be  used  for  acetylene  be- 
cause the  flame  produces  soot.  A  common  form  of  burner 
is  shown  in  Fig.  48.  The  acetylene  as  it  escapes  from  the 
supply  pipe  (A)  into  the  burner  sucks  in  air  through  the 
small  side  holes  (B).  This  mixture,  upon  ignition,  burns 
as  a  small  flat  flame  at  right  angles  to  the  burner  (Fig.  49). 
(See  Part  II,  Exp.  87.) 


Fig.  48.  —  Acetylene  Burner.  Fig.  49.  —  Acetylene  Flame. 

194.  Generation  of  Acetylene.  —  The  gas  can  be  generated  by 
putting  the  calcium  carbide  into  a  flask  provided  with  a  dropping 
funnel  and  delivery  tube,  and  allowing  water  to  drop  slowly  upon 
the  carbide;  the  gas  thus  generated  can  be  collected  in  bottles 
over  water.      There  are  two  classes  of  commercial  generators.      In 
small  ones  water  is  added  to  the  calcium  carbide;  such  generators 
may  be  safely  used  on  the  lecture  table  and  in  a  bicycle  or  auto- 
mobile lantern.      In  larger  generators  calcium  carbide  drops  auto- 
matically into  a  large  volume  of  water  and   the  gas  collects  in  a 
reservoir  which  by  its  pressure  regulates  the  machine.      A  solution 
of    acetylene  in  acetone    is   often  used  as  a  source  of  the  gas, 
especially  for  automobiles. 

195.  Oxy-acetylene  Flame.  —  A  mixture  of  acetylene 
and   considerable    oxygen  burns  with  an  intensely  hot 
flame,  the  temperature  being  nearly  3000°  C.     The  oxy- 
acetylene  flame  is  used  to  sever  large  pieces  of  metal. 
Ordinary   tools   cut  hard   metals   slowly,   but   the   oxy- 
acetylene  flame  when  its  tip  is  passed  slowly  across  the 


174  CHEMISTRY 

metal  burns  its  way  through  steel  shafts  and  girders 
with  astonishing  rapidity.  Metal  structures,  such  as 
fences,  bridges,  frames  of  buildings,  abandoned  battle- 
ships, etc.,  are  speedily  dismantled  by  this  flame.  The 
fire  department  of  large  cities  is  equipped  with  an  oxy- 
acetylene  outfit  for  cutting  a  passage  through  steel  doors 
of  vaults  or  effecting  an  entrance  into  parts  of  a  fire- 
proof building.  The  oxy-acetylene  flame  is  also  used  for 
welding  metals.  The  burner  (Fig.  50)  is  constructed  like 

^^ the  oxy-hydrogen  blowpipe  (26). 

"^|  Oxygen  enters  at  A  and  acetylene 

^  at  B.    The  tip  of  the  burner  does 

not  melt  because  the  flame  at  that 

Fig.  50.  —  Oxy-acetylene  ...  , 

Blowpipe.  P°mt  1S  not  very  hot 

196.   Petroleum  is  the  source  of 

many  useful  hydrocarbons.  It  is  obtained  from  the  earth 
in  many  parts  of  the  world.  In  the  United  States  the 
chief  localities  are  Ohio,  Pennsylvania,  Texas,  and  Cali- 
fornia. The  immense  deposits  in  Russia  are  in  the  Baku 
district  on  the  Caspian  Sea. 

Crude  petroleum  is  an  oily  liquid  with  an  unpleasant 
odor.  Its  color  varies  from  amber  to  black;  some  kinds 
are  greenish  in  reflected  light.  Its  composition  is  com- 
plex, but  all  varieties  are  essentially  mixtures  of  liquid 
hydrocarbons  in  which  gaseous  and  solid  hydrocarbons 
are  dissolved.  Some  varieties  contain  compounds  of 
nitrogen  and  of  sulphur. 

In  some  localities  the  oil  issues  directly  from  the  earth,  but  it  is 
usually  necessary  to  drill  a  deep  hole  and  insert  a  pipe  into  the  porous 
rock  containing  oil.  Sometimes  the  oil  "  spouts  "  out  of  the  well  when 
first  drilled,  but  usually  a  pump  is  needed  to  draw  it  to  the  surface. 

Some  crude  petroleum  is  used  in  making  water  gas  and 
as  fuel  for  locomotives  and  steamships,  but  most  of  it  is 


CARBON  175 

separated  into  various  commercial  products.  This  process 
is  called  refining.  The  petroleum  is  distilled  by  heating 
it  in  huge  iron  vessels,  called  stills  or  retorts,  and  condens- 
ing the  vapors  in  coiled  pipes  immersed  in  cold  water. 
This  process  is  sometimes  called  fractional  distillation, 
since  the  petroleum  is  merely  separated  into  parts  or  frac- 
tions. (Compare  176,  second  ^f.)  Certain  products  are 
obtained  from  the  residue  left  in  the  still. 

The  portions  that  pass  over  between  certain  temperatures  are 
collected  in  separate  tanks,  and  further  separated  and  purified  by 
redistillation.  The  chief  commercial  products  thus  obtained  are  pe- 
troleum ether  (boiling  point  4o°-7o°  C),  gasoline  (ljo°-go°),  naphtha 
(90°-!  20°),  benzine  (i2o°-i5o°),  kerosene  (i5o°-3oo°);  there  are 
various  commercial  grades  of  these  products,  which  are  distinguished 
by  the  range  of  boiling  points  or  by  the  specific  gravity.  These 
liquids  are  mixtures  of  several  different  hydrocarbons. 

The  petroleum  products  enumerated  above  are  extensively  used 
as  solvents,  fuels,  and  illuminants.  At  present  gasoline  is  the  most 
suitable  fuel  for  the  internal  combustion  engines  used  in  propelling 
automobiles,  motor-cycles,  motor  boats,  and  flying  machines.  The 
vapor  of  gasoline  burns  readily.  If  the  vapor  is  mixed  with  air 
and  the  mixture  is  ignited  by  an  electric  spark,  the  combustion  is 
so  rapid  that  it  is  practically  an  explosion;  the  suddenly  expanded 
gases  exert  pressure,  which  is  converted  by  the  machinery  into 
steady  and  continuous  motion.  Kerosene  is  so  widely  used  as  an 
illuminant  that  it  is  carefully  freed  from  readily  inflammable  liquids 
and  gases,  which  might  cause  an  explosion,  and  from  tarry  matter 
and  semi-solid  hydrocarbons,  which  would  clog  a  lamp  wick.  This 
is  done  by  agitating  it  successively  with  sulphuric  acid,  sodium 
hydroxide,  and  water.  Commercial  kerosene  must  have  a  legal 
flashing  point.  This  is  "  the  temperature  at  which  the  oil  gives 
off  sufficient  vapor  to  form  a  momentary  flash  when  a  small  flame 
is  brought  near  its  surface."  The  legal  minimum  flashing  point  in 
most  states  is  about  110°  Fahrenheit  (about  44°  C.). 

From  the  residue  (left  in  the  still  after  the  distillation  of  the 
kerosene)  many  grades  of  lubricating  oil  are  obtained;  also  vaseline 
and  paraffin.  Mineral  lubricating  oils  have  largely  replaced  animal 


i  y6  CHEMISTRY 

and  vegetable  oils.  Vaseline  finds  extensive  use  as  an  ointment. 
Paraffin  wax  is  made  into  candles  and  into  a  water-proof  coating 
for  many  substances.  The  final  residue  in  the  still  is  mainly 
carbon  and  is  called  petroleum  coke;  it  is  made  into  electric  light 
carbons.  The  oils  separated  from  petroleum  are  often  called  mineral 
oils,  because  they  differ  in  composition  from  oils  obtained  from 
plants  and  animals.  (See  Fats,  Chapter  XVII.) 

197.  Coal  Tar  is  another  source  of  hydrocarbons.    This 
is  a  viscous  liquid  obtained  as  a  by-product  in  the  manu- 
facture of  coal  gas  (205).    Like  petroleum  it  is  a  complex 
mixture,  and  by  its  distillation  are  obtained  the  hydro- 
carbons benzene  (also  called  benzol,  C6H6),  naphthalene 
(CioHg) ,  and  anthracene  (CnHio),  as  well  as  many  other  car- 
bon compounds,  especially  phenol.     (See  Part  II,  Exp.  99.) 

198.  Benzene  is  a  light,  colorless  liquid.    It  burns  with 
a   luminous,    smoky   flame.      Ordinary   illuminating   gas 
owes  its  luminosity  partly  to  benzene  (208).    It  dissolves 
fats,  resins,  iodine,  sulphur,  and  rubber.     It  should  not 
be  confused  with  benzine,  which  is  a  mixture  of  other 
hydrocarbons    obtained    from    petroleum.      Benzene    is 
chiefly  used  in  preparing  organic  compounds,  e.g.  nitro- 
benzene (CeHs-NQa),  a  yellow  poisonous  liquid  which  has 
the  odor  of  bitter  almonds.     Nitrobenzene  is  made  into 
aniline  (C6H5.NH2)  from  which  are  manufactured  aniline 
dyes.    Naphthalene  (CioHs)  is  a  white,  lustrous  crystalline 
solid  having  a  penetrating,  disagreeable  odor.    It  is  familiar 
under  the  name  of  "moth  balls"  and  is  used  as  a  substitute 
for  camphor  in  protecting  cloth  from  moths.     From  an- 
thracene (Ci4Hio)  is  made  alizarin,  which  is  a  dyestuff 
used  to  produce  brilliant,  fast  colors.    Phenol  (C6H5.OH) 
is  a  white  crystalline  solid.    It  has  a  peculiar,  smoky  odor. 
A  solution  of  phenol  in  water  is  called  carbolic  acid  and  is 
used  as  a  disinfectant. 


CARBON  177 

Carbides 

199.  Carbides  are  compounds  of  carbon  and  certain 
elements,  especially  metals.     Many  are  made  by  subject- 
ing a  mixture  of  carbon  and  oxides  to  the  intense  heat 
generated  in  an  electric  furnace. 

200.  Silicon   Carbide    or   Carborundum,    SiC. --This 
carbide  is  a  crystallized  solid,  which  varies  in  color  from 
white  to  brownish-green  or  black.     It  is  extremely  hard, 
being  nearly  as  hard  as  diamond.    Acids  do  not  affect  it, 
but  it  is  decomposed  by  fusion  with  potassium  hydroxide 
or  other  alkalies.    The  extreme  hardness  of  carborundum 
has  led  to  its  extensive  application  as  an  abrasive,  and 
large  quantities  are  made  into  a  great  variety  of  grinding 
wheels,  whetstones,  and  polishing  cloths.     (See  Part  II, 
Exp.  93.) 

Carborundum  is  manufactured  by  fusing  a  mixture  of  sand 
(silicon  dioxide,  SiO2)  and  coke  (carbon,  C)  in  an  electric  furnace 
constructed  on  the  resistance  type  (Fig.  51).  It  is  essentially  a  large 
rectangular  box  with  per- 
manent ends  and  loosely 
built  sides.  Each  end  is 
provided  with  a  heavy 
metal  plate.  The  wires 
for  the  electric  current  Fig  5I> —Electric  Furnace  —  Resistance  Type, 
are  attached  to  the  outer 

ends  of  these  plates,  while  the  huge  carbon  electrodes  fit  into  the 
inner  ends,  and  project  into  the  furnace.  A  cylindrical  core  of 
granulated  coke  makes  an  electrical  connection  between  the  elec- 
trodes. The  mixture  of  sand  and  coke  (to  which  salt  and  sawdust 
are  added  to  contribute  to  the  fusion  and  porosity)  is  packed  around 
this  core  inside  the  box.  The  heat  generated  by  the  resistance  of 
the  carbon  core  to  the  passage  of  the  powerful  current  of  electricity 
produces  a  chemical  change  essentially  as  follows:  - 

Si02      +      3C      =    SiC       +        2CO 

Silicon  Dioxide         Carbon      Silicon  Carbide      Carbon  Monoxide 


i78 


CHEMISTRY 


When   the  operation  is  over,  the  furnace  is  allowed  to  cool,  the 
side  walls  are  pulled  down,  and  the  carborundum  is  removed. 

201.  Calcium  Carbide,  CaC2,  is  a  brittle,  dark  gray, 
crystalline  solid.  The  most  striking  and  useful  property 
is  its  action  with  water,  acetylene  being  formed,  thus:  - 


CaC2 

Calcium  Carbide 


2H2O     = 

Water 


C2H2 

Acetylene 


Ca(OH)2 

Calcium  Hydroxide 


Calcium  carbide  is  used  to  generate  acetylene  gas.  This 
gas  burns  with  a  brilliant  flame,  and  is  used  as  an  illumi- 
nant  (193,  194). 

Calcium  carbide  is  made  from  lime  (calcium  oxide,  CaO)  and 
coke  or  coal  in  an  electric  furnace.  The  chemical  change,  like  that 
in  the  manufacture  of  carborundum,  is  caused  solely  by  the  in- 
tense heat  and  may  be  represented  thus:  — 

3C    +       CaO       =        CaC2       +          CO 

Carbon        Calcium  Oxide       Calcium  Carbide       Carbon  Monoxide 


The  furnace  in  which 


Fig.  52.  —  Electric  Furnace  for  Making 
Calcium  Carbide. 


calcium  carbide  is  made  is  sketched  in 
Fig.  52.  The  mixture  of 
coke  and  lime  (shown  in 
the  furnace)  is  intro- 
duced through  the  trap 
cover  A  and  slowly  sinks 
down  into  the  space  where 
the  intense  heat  is  pro- 
duced by  the  electricity 
as  it  passes  between  the 
electrodes  G  and  E,  E. 
The  liquid  calcium  car- 
bide is  drawn  off  through 
F.  The  carbon  monoxide 
rises  through  the  pipes 
D,  D  and  enters  the  up- 
per part  of  the  furnace, 
together  with  air  sup- 
plied through  C,  C;  this 


CARBON  179 

mixture  burns  and  heats  the  coke  and  lime.    The  waste  gases  (carbon 
dioxide  and  nitrogen)  escape  through  B. 

Cyanogen  and  Related  Compounds 

202.  Cyanogen,  C2N2  (sometimes  written  (CN)2),  is 
a  colorless  gas  with  the  odor  of  peach  kernels.  It  is  ex- 
ceedingly poisonous.  It  burns  with  a  purplish  flame.  It 
may  be  prepared  by  heating  mercuric  cyanide  (Hg(CN)2). 
Cyanogen  is  a  radical,  and  in  compounds  it  acts  like  an 
element.  Its  corresponding  acid  is  hydrocyanic  or  prus- 
sic  acid  (HCN).  This  acid  is  prepared  by  heating  a 
cyanide  with  sulphuric  acid,  just  as  hydrochloric  acid  is 
obtained  from  a  chloride.  It  is  a  colorless  volatile  liquid. 
The  vapor  is  sometimes  used  in  treating  trees  affected  with 
San  Jose  scale.  The  solution  smells  like  peach  kernels. 
Both  vapor  and  solution  are  exceedingly  poisonous. 
Potassium  cyanide  (KCN)  and  sodium  cyanide  (NaCN) 
are  salts  of  hydrocyanic  acid.  They  are  white,  deli- 
quescent solids.  Both  are  deadly  poisons;  they  should 
not  be  touched  with  the  hands,  and  unusual  care  must  be 
taken  not  to  inhale  small  particles  or  even  the  gases  that 
escape  from  bottles  containing  cyanides.  They  are  used 
in  gold  and  silver  plating  and  in  the  cyanide  process  of 
extracting  gold  from  its  ores. 

EXERCISES 

1.  In  what  forms  does  free  carbon  occur  in  nature?    Name  ten  familiar 
solids,  three  liquids,  and  two  gases  which  contain  carbon. 

2.  What  is  the  chemical  relation  of  graphite  to  diamond,  and  how  can 
this  relation  be  proved?     State  the  source,  properties,  and  uses  of  graphite. 

3.  State  the  properties  and  uses  of   (a)  wood  charcoal  and  (b]  animal 
charcoa  .    Give  a  brief  account  of  both  methods  of  preparing  wood  charcoal. 
Charcoal  when  burned  often  leaves  a  white  residue;  why? 

4.  Write  an  essay  of  one  hundred  words  on  one  or  more  oc  these  topics: 
(a)  Famous  diamonds.  (6)  Coal  as  fuel,  (c)  Carbon  in  electrical  industries. 


i8o  CHEMISTRY 

(d)  Carbon  in  paint  making,  (e)  The  cycle  of  carbon  in  nature.  (/)  History 
of  carbon  dioxide. 

5.  Describe  fully  the  action  of  carbon  dioxide  on  calcium  hydroxide. 
Express  the  reaction  by  an  equation. 

6.  What  is  the  test  for  (a)  carbon,  (b}  carbon  monoxide,  (c}  carbon 
dioxide,  (d)  a  carbonate? 

7.  Suggest  methods  of  proving  that  (a)  CO2  is  the  formula  of  carbon 
dioxide  and  (b\  CO  of  carbon  monoxide. 

8.  Give  the  equations  for  (a)  the  oxidation  of  carbon  to  carbon  monox- 
ide and  (V)  the  reduction  of  carbon  dioxide  to  carbon  monoxide. 

9.  Illuminating  gas,  water  gas,  and  the  gas  that  escapes  from  a  coal 
fire  are  poisonous.     Why?    What  is  a  pulmotor  and  for  what  is  it  used? 

10.  Describe  acetylene.    How  is  it  prepared?    Give  the  equation  for  the 
reaction.     Summarize  the  properties  of  acetylene. 

11.  Describe    the   acetylene  (a)  flame,  (b)  burner,  and  (c)  generator. 

12.  Essay  topics:    (a)  The  oxygen  helmet  and    its  use.     (b)  Miners' 
safety  lamps.     (See  Miners'  Circulars  4  and  12,  Bureau  of  Mines,  Wash- 
ington, D.C.) 

13.  What  is  kerosene?     Define  the  term  flashing  point. 

14.  What  properties  of  amorphous  carbon,  e.g.  lampblack,  make  it  a 
useful  ingredient  of  printer's  ink  and  of  paint? 

PROBLEMS 

1.  Calculate  the  percent  of  carbon  in  (a)  carbonic  anhydride,  (6)  acid 
calcium  carbonate,  (c)  carbonic  oxide,  (d)  C^H^On,  (e)  C2H6O,  (/)  marble. 

2.  What  weight  of  carbon  is  contained  in  2  1.  of  carbon  monoxide?     In 
2  1.  of  carbon  dioxide?     (Standard  conditions.) 

3.  What  volume  at  standard  conditions  is  occupied  by  (a)  70  gm.  of  car- 
bon dioxide,  and  by  (b)  45  gm.  of  carbon  monoxide? 

4.  The  volume  of  one  gram  of  carbon  dioxide  is  measured  at  20°  C. 
and  765  mm.     What  is  its  volume? 

5.  If  12  grams  of  carbon  are  burned   to  carbon  dioxide,  what  will  be 
the   volume  of  the  gas  compared  with  i  gram  of  hydrogen  at  the  same 
temperature  and  pressure? 

6.  How  much  oxygen  by  weight  and  by  volume  (standard  conditions) 
is  needed  to  convert  the  following  into  carbon  dioxide:    (a)   i  kg.  of  pure 
charcoal,    (b)   a    diamond    weighing   250  milligrams,  and    (c)    30  gm.    of 
graphite? 

7.  A  volume  of  carbon  dioxide  measures  575  cc.  at  16°  C.  and  765  mm. 
What  will  be  its  volume  and  weight  under  standard  conditions? 

8.  Thirty  grams  of  carbon  are  burned  to  carbon  dioxide.    What  weight 
of  potassium  chlorate  must  be  decomposed  to  provide  the  oxygen? 


CHAPTER  XVI 
ILLUMINATING  GASES  — FLAMES 

203.  Illuminating    Gases. — Besides    acetylene    (193) 
there  are  other  kinds  of  illuminating  gas.     Pintsch  gas 
is  made  from  petroleum  by  heating  the  vaporized  oil  to 
a  high  temperature  and  is  often  called  an  oil  gas.     It 
burns  with  a  brilliant  flame  and  is  used  for  lighting  rail- 
way cars.    Coal  gas  and  water  gas,  however,  are  the  most 
important. 

204.  Coal  Gas  is  made  by  heating  bituminous  coal  out 
of  contact  with  air  and  purifying  the  volatile  product. 
(See  Part  II,  Exps.  94,  98,  99, 100.)     (Compare  176,  196.) 

205.  Manufacture  of  Coal  Gas.  —  A  diagram  of  a  coal  gas  plant 
is  shown  in  Fig.  53.      The  coal  is  heated  several  hours  in  closed 
retorts.      The   volatile    products   pass   from    the   retorts   into   the 
hydraulic  main.      Here  some  of  the  tar  is  deposited  and  ammonium 
compounds   are   dissolved    by   the   water   which  flows   constantly 
through  the  main.      The  ammoniacal  liquor  and  tar  flow  into   the 
tar  well.      From  the  hydraulic  main  the  hot  and  impure  gas  passes 
through  the  condenser.      The  main  object  of  the  condenser  is   to 
cool  the  gas  and  remove  tar.      An  exhauster  transfers  the  gas  from 
the  condenser  into  the  scrubber  (and  onward  through  the  purifiers 
into   the  gas  holder).      The  purpose  of  the  scrubber  is  to  remove 
the    remaining    ammonium    compounds,   the    carbon  dioxide    and 
hydrogen  sulphide,  and   the  last  traces  of  tar.      From  the  scrubber 
the  gas  passes  into  the  purifiers.      These  contain  lime  or  iron  oxide, 
or  both,  which  remove  any  remaining  carbon  dioxide  and  sulphur 
compounds.      The  purified  gas  next  passes  through  a  large  meter, 
which  records  its  volume,  into  a  gas  holder  from  which  the  gas  is 
forced  through  the  pipes  to  the  consumer. 


182 


CHEMISTRY 


A  ton  of  good  gas  coal 
yields  about  10,000  cubic  feet 
of  gas,  1400  pounds  of  coke" 
(178),  120  pounds  of  tar,  20 
gallons  of  ammoniacal  liquor 
(101),  and  a  varying  amount 
of  gas  carbon  (179).  The 
tar,  or  coal  tar  as  it  is  often 
called,  collected  from  the  hy- 
draulic main  and  condenser, 
is  a  thick,  black,  foul  smell- 
ing liquid.  Some  is  used  for 
preserving  timber,  making 
tarred  paper  and  black  var- 
nishes, and  as  a  protective 
paint.  Most  of  it  is  sepa- 
rated by  distillation  into  its 
important  constituents  (197). 

206.  Water  gas    is 

made  by  forcing  steam 
through  a  mass  of  hot 
coke  or  anthracite  coal 
and  mixing  the  gaseous 
product  with  hot  gases 
obtained  from  oil.  It  is 
essentially  a  mixture  of 
hydrogen,  carbon  mon- 
oxide, and  a  small  pro- 
portion of  hydrocarbons. 

207.  Manufacture    of 
Water  Gas. —  The   essential 
parts  of  the  apparatus    are 
shown    diagrammatically  in 
Fig.  54.      Air  is  forced  by  a 
blower  through    the   fire   in 


ILLUMINATING   GASES 


183 


02 


^  i$^$^$^^^^^§$$^s$^^$s$$?^ 


O 


I  he  generator,  and  the  hot 
gases  that  are  produced 
pass  down  the  carbureter, 
up  into  the  superheater, 
and  escape  through  an 
opening  (not  shown)  into 
the  open  air.  This  opera- 
tion heats  the  fire  brick 
inside  the  carbureter  and 
superheater  intensely  hot, 
air  often  being  forced  in  to 
raise  the  temperature.  The 
air  valves  and  the  opening 
at  the  top  of  the  super- 
heater are  now  closed,  and 
steam  is  forced  into  the 
^  generator  at  the  bottom. 
"£  In  passing  through  the 
J  mass  of  incandescent  car- 
2  bon,  the  steam  and  carbon 

2  interact  thus:  — 

I 

3  C  +  H2O  =  CO  +    H2 

I    Carbon    Steam   Carbon     Hydrogen 
4-  Monoxide 


•^  The  mixed  gases  rise  to 
the  top '  of  the  carbureter, 
where  they  meet  a  spray 
of  oil,  and  as  the  gaseous 
mixture  passes  down  the 
carbureter  and  up  the 
superheater,  the  hydro- 
carbons of  the  oil  are 
transformed  by  the  intense 
heat  into  gaseous  hydro- 
carbons which  do  not 
liquefy  when  the  final  gas 
is  cooled.  From  the  super- 
heater the  water  gas  passes 


1 84 


CHEMISTRY 


through  the  purifying  apparatus  into  a  holder.  The  oil  is  added  to 
provide  illuminants,  since  the  flame  from  the  hydrogen- carbon  mon- 
oxide mixture  is  very  feeble. 

Water  gas  is  seldom  burned  alone,  but  is  usually  mixed 
with  60  or  70  per  cent  of  coal  gas.  This  mixture  is  pop- 
ularly called  "illuminating  gas."  Owing  to  the  high 
percentage  of  carbon  monoxide,  water  gas  and  mixtures 
containing  it  are  poisonous  (189). 

208.  Characteristics  of  Illuminating  Gases.  —  Illu- 
minating gases  have  a  disagreeable  odor.  They  are 
complex  mixtures.  The  following  table  shows  the  ap- 
proximate composition  of  average  samples :  - 

COMPOSITION  OF  ILLUMINATING  GASES   (By  VOLUME) 


Constituents 

Coal  Gas 

Water  Gas 

Oil  Gas 

Methane  .          

-2  A  .  C 

10-8 

188 

Illuminants 

C  O 

16  6 

4.e  o 

Hydrogen  

40.  0 

32.1 

Carbon  Monoxide 

7  2 

26  i 

Carbon  Dioxide  

I.I 

3.0 

Nitrogen 

•2  2 

2  4 

I  I 

Methane,  hydrogen,  and  carbon  monoxide  burn  with  a  feeble 
(non-yellow)  flame,  and  are  often  called  diluents;  they  furnish  heat, 
but  no  light.  The  illuminants  consist  of  ethylene  (C2H4),  acetylene 
(C2H2),  benzene  (C6H6),  and  other  hydrocarbons. 

The  luminosity  of  an  illuminating  gas  is  measured  by 
a  photometer  and  is  expressed  in  candles  or  candle  power. 
The  determination  is  made  by  comparing  the  light  pro- 
duced by  burning  the  gas  with  the  light  produced  by 
a  standard  wax  candle  or  a  standard  flame.  The  candle 
power  of  coal  gas  is  about  17,  of  water  gas  about  25,  and 


ILLUMINATING   GASES 


185 


of  oil  gas  50  or  more.  Ordinary  illuminating  gas  has  a 
varying  candle  power,  since  it  is  usually  a  mixture  of 
coal  gas  and  water  gas.  The  use  of  mantles  has  greatly 
improved  the  methods  of  lighting  by  gas  (213). 


Flames 

209.  General  Nature  of  Flames.  --  The  term  flame 
is  ordinarily  applied  to  a  light  produced  by  burning  gases 
in  air.  If  the  flame  contains  much  free  carbon,  the  un- 
consumed  particles  are  rendered  more  or  less  incandescent 
by  the  heat  liberated  by  the  chemical  change  and  the 
flame  is  luminous.  It  should  not  be  overlooked  that  the 
flame  from  burning  liquids  and  solids  is  due  to  burning 
gases.  In  an  illuminating  gas  flame  the  gas  itself,  of 
course,  is  burning  in  air.  In  a  lamp  flame  the  gas  that 
burns  comes  from  the  oil  that  is  drawn 
up  the  wick  by  capillary  attraction,  and 
then  volatilized  by  the  heat.  Similarly^ 
in  a  candle  flame  the  burning  gas  comes 
from  the  melted  and  volatilized  wax. 

Ordinary  flames  are  due  primarily  to  the 
combination  of  the  oxygen  of  the  air  with 
the  gas  or  its  elementary  constituents. 
We  usually  burn  gases  in  an  abundance 
of  air,  but  a  flame  is  produced,  though 
not  so  conveniently,  if  the  conditions  are 
reversed. 


Fig.  55-  —  Com- 
bustion in  Illu- 
minating Gas 
and  in  Air. 


A  simple  experiment  illustrates  this  fact.  In 
the  apparatus  shown  in  Fig.  55,  the  lamp  chimney 
B  is  filled  with  illuminating  gas  through  the  bent 
tube  D,  and  its  escape  is  temporarily  prevented  by  closing  the  open- 
ing in  the  asbestos  cover  A.  The  gas  is  lighted  at  the  lower  end 
of  the  tube  C,  and  when  the  hole  in  A  is  uncovered,  the  flame  rises 


i86 


CHEMISTRY 


in  C  and  continues  at  the  end  within  the  chimney  as  long  as  air  is 
drawn  up  through  C  and  gas  supplied  through  D.  The  unconsumed 
illuminating  gas  escapes  through  the  hole  in  A,  and  if  ignited,  burns 
much  like  the  other  flame,  as  shown  in  the  figure.  The  gas  pres- 
sure must  be  carefully  regulated  to  produce  a  satisfactory  result. 
Chemically  both  flames  are  alike.  The  outer  flame  is  in  an  atmos- 
phere of  air,  while  the  inner  flame  is  in  an  atmosphere  of  illuminat- 
ing, gas;  but  both  flames  are  due  to  the  combination  of  oxygen 
with  the  elementary  constituents  of  the  illuminating  gas. 

210.  Structure  and  Characteristics  of  Luminous  Flames. 

-  The  luminous  hydrocarbon  flame  has  several  distinct 
parts,  and  the  structure  of  the  flame  is 
essentially  the  same,  whether  produced  by 
burning  illuminating  gas,  kerosene  oil,  or 
candle  wax.  The  candle  flame  may  be 
taken  as  the  type.  Examination  of  the 
sketch  of  an  enlarged  vertical  section 
shown  in  Fig.  56  reveals  four  somewhat 
conical  portions,  (i)  Around  the  wick 
there  is  a  dark  cone  (A),  filled  with  com- 
bustible, but  unburned,  gases  formed  from 
the  melted  wax.  As  already  stated  it  is 
possible  to  draw  off  these  gases  and  light 
them.  (2)  Around  the  lower  part  of  the 
dark  cone  is  a  faint  bluish  cup-shaped  part 
(B,  B).  It  is  the  lower  portion  of  the  exterior  cone  where 
complete  combustion  of  the  gases  occurs,  since  plenty  of 
oxygen  from  the  air  reaches  this  portion.  (3)  Above  the 
dark  cone  is  the  luminous  portion  (C).  It  is  the  largest 
and  most  important  part  of  the  flame.  It  is  usually 
spoken  of  as  "the  flame."  Combustion  is  incomplete 
here,  because  little  or  no  oxygen  can  pass  through  the 
exterior  cone.  The  temperature  is  high,  however,  and 
the  hydrocarbons  undergo  complex  changes.  Acetylene 


Fig.  56.  —  Parts 
of  a  Typical 
Candle  Flame. 


ILLUMINATING   GASES  187 

is  probably  formed.      The  most  characteristic  change  is 

the  liberation  of  small  particles  of  carbon.    This  liberated 

carbon,  heated  to  incandescence  by  the  burning  gases, 

makes  the  flame  luminous.     (4)  The  exterior  cone  (D,  D) 

is  almost  invisible.     Here    the   combustion  is  complete, 

because  the  oxygen  of  the  air  changes 

all  the  carbon  into  carbon  dioxide.  That 

this  is  the  hottest  region  of  the  flame 

can  be  shown   by  pressing  a  piece  of 

stiff  white  paper  for  an  instant  down 

upon  the  flame  almost  to  the  wick.    The 

paper  will  be  charred  by  the  hot  outer    Fig    57._  charred 

portion  of  the  flame,  as  shown  in  Fig.       Paper  Showing  the 

57.     (See  Part  II,  Exp.  95.)  Hottest  Part  of  a 

e  ^.  i       r         j    •  Candle  Flame. 

These  four  portions  may  be  found  in 
all  luminous  hydrocarbon  flames,  whatever  the  shape.  An 
ordinary  gas  flame  is  flattened  by  forcing  the  gas  through 
a  narrow  slit  in  the  burner  tip,  hence  the  flame  gives  more 
light. 

The  gaseous  products  of  the  combustion  of  hydrocar- 
bons are  water  vapor  and  carbon  dioxide.  A  bottle  in 
which  a  candle  is  burning  has,  at  first,  a  deposit  of  moisture 
on  the  inside;  and  if  the  candle  is  removed  and  calcium 
hydroxide  solution  added,  the  presence  of  carbon  dioxide 
is  shown  by  the  cloudiness  of  the  solution.  The  oxygen 
needed  by  the  burning  hydrocarbons  is  obtained  from  the 
air.  If  not  enough  oxygen  is  present,  the  flame  smokes, 
i.e.  the  carbon  is  thrown  off  into  the  air  before  the  par- 
ticles are  heated  hot  enough  to  glow.  All  oil  lamps  are 
so  constructed  that  air  enters  the  burner  below  the  flame. 

The  luminosity  of  hydrocarbon  flames  is  affected  by  tempera- 
ture. A  candle  flame  may  be  cooled  enough  to  extinguish  it. 
Thus,  if  a  coil  of  copper  wire  is  lowered  upon  a  candle  flame,  the 


i88 


CHEMISTRY 


flame  smokes,  loses  its  yellow  color,  and  finally  goes  out;  but  if  a 
coil  of  hot  wire  is  used,  the  flame  burns  unchanged  (Fig.  58). 

Not  all  luminous  flames  are  hydrocarbon  flames.      Thus,  mag- 
nesium burns  with  a  brilliant  flame.      Its  luminosity  is  due  to  the 
incandescence  of  solid  par- 
ticles of  magnesium  oxide. 
Similarly,  the  bright  flame 
of  burning  phosphorus  is 
accounted  for   by  the  in- 
candescent    particles      of 
solid   phosphorus   pentox- 
ide.      (See  also  213.) 


Fig.  58.  —  Effect  of  Lowering  the  Tempera- 
ture of  a  Candle  Flame. 


o 


211.   Non-Lumi- 
nous   Flames.  --  The 

hydrogen  flame  is  almost  invisible  in  air  and  oxygen, 
but  pale  blue  in  chlorine.  The  flames  of  car- 
bon monoxide  and  methane  are  also  faint  blue. 
The  most  common  non-luminous  flame  is  the 
Bunsen  flame. 

212.   The  Bunsen  Burner  and  its  Flame.  - 
When  illuminating  gas  is  mixed  with  air  before 
burning,  and  the  mixture  burned  in  a  suitable 
burner,  a  flame  is  produced  which  is  non-lumi- 
nous, very  hot,  and  deposits  no  carbon.    Such  a 
flame  is  called  the  Bunsen  flame,  for  it  was  first 
produced  in  a  burner  devised  by  the  German 
chemist  Bunsen.      This  burner  is  used  in  lab- 
oratories as  a  source  of  heat  (Fig.  59).      The 
gas  enters  the  base  and  escapes  through  a  very 
small  opening  into  the  long  tube, 
which  screws  down  over  this  open- 
ing.   At  the  lower  end  of  the  long 
(     ™    .     tube  there  are  two  holes,  through 

Fig.  59.  —  Parts  of  a  Typi- 
cal Bunsen  Burner.          which  air  is  drawn  by  the  gas  as 


ILLUMINATING   GASES  189 

it  rushes  out  of  the  small  opening.  The  gas  and  air 
mix  as  they  rise  in  the  tube,  and  this  mixture  of  air 
and  gas  burns  at  the  top  of  the  long  tube.  The  size 
of  the  air  holes  at  the  bottom  of  the  long  tube  may 
be  changed  by  a  movable  ring,  thus  varying  the  vol- 
ume of  the  entering  air.  When  the  holes  are  open,  the 
typical  non-luminous,  hot  Bunsen  flame  is  formed.  The 
combustion  of  the  constituents  of  the  hydrocarbons  is 
practically  complete.  The  non-luminous  flame  is  free 
from  soot,  therefore  apparatus  heated  by  this  flame  is  not 
blackened. 

The  gas  burns  at  the  top  of  the  tube  and  not  inside, 
because  the  proper  mixture  of  gas  and  air  flows  out  more 
quickly  than  the  flame  can  travel  back  through  the  tube 
to  the  small  exit.  If  the  gas  supply  is  slowly  decreased, 
the  flame  becomes  smaller,  disappears  with  a  slight  ex- 
plosion, and  burns  at  the  small  gas  opening  inside  the 
tube.  A  sudden  draft  of  air,  too  large  holes  at  the  lower 
end  of  the  tube,  or  too  low  gas  pressure  also  may  cause 
the  flame  to  "  strike  back,"  as  this  action  is  called.  This 
change  is  due  to  the  fact  that  the  tube  contains  an  explo- 
sive mixture  of  air  and  illuminating  gas,  through  which  the 
flame  travels  downward  faster  than  the  mixture  escapes 
from  the  tube.  This  modified  flame,  which  has  a  pale 
color,  a  disagreeable  odor,  and  deposits  soot,  should  be 
extinguished  and  the  proper  flame  produced  before  further 
use. 

The  Bunsen  flame  has  many  characteristic  properties. 
Its  color  is  bluish,  and  the  cones  have  different  tints. 
The  outermost  cone  is  not  easily  distinguished;  so  for 
practical  purposes  it  is  convenient  to  divide  the  flame 
into  two  parts,  —  an  inner  cone  of  unburned  gases  and 
an  outer  cone  in  which  all  the  carbon  is  consumed. 


190 


CHEMISTRY 


The  existence  of  these  two  cones  can  be  shown  by  simple  ex- 
periments.     A  match   held  near  the  outer  part  of  the  flame  takes 
fire  quickly.      Combustible  gases  can  be  drawn  off  by  a  tube  from 
the  inner  cone  and  ignited,  as  in  the  candle  flame.      A  match  laid 
for  an  instant  across  the  top  of  the  tube  is  charred  only  at  the 
two  points  where  it  touches  the  outer  cone;  and  a 
fi  sulphur  match  suspended  by  a  pin  across  the  top 

of  an  unlighted  burner  is  not  kindled  until  some 
~~  |€LDr~°  time  after  the  gas  is  first  lighted  (Fig.  60).  Finally, 
a  wire  gauze,  if  pressed  down  upon  the  flame,  shows 
a  dark  central  portion  surrounded  by  a  luminous 
ring  due  to  the  inner  and  outer  cones  respectively 
(see  Fig.  47).  (See  Part  II,  Exp.  96.) 


213.   Welsbach    Light.  -  -  The    Bunsen 


-A 


-B 


Fig.     60.  —  Ex- 
periment Show- 

ing  the  Inner  flame  1S  extensively  used  in  producing  the 
Cone  of  Un-  Welsbach  light.  The  non-luminous  flame 
burned  Gases  neats  a  mantle  consisting  of  a  firm  network 

in   the  Bunsen      -  -    ,, 

Flame  99  Per  ce  thorium  ox- 

ide and  i  per  cent  of  cerium 
oxide,  and  the  mantle  glows  with  an  intense 
light.  The  candle  power  varies  from  40  to 
100.  (See  Part  II,  Exp.  102.) 

214.  Oxidizing  and  Reducing  Flames.  - 
The  outer  portion  of  the  Bunsen  flame  is 
called  the  oxidizing  flame,  because  here  oxy- 
gen is  abundant.  The  inner  portion  is  called 
the  reducing  flame,  because  here  the  excess 
of  reducing  gases  withdraws  oxygen  from  ox- 

Fig.  oi. — The 

ides.  A  sketch  of  the  general  relation  of  Oxidizing 
these  flames  is  shown  in  Fig.  61.  A  is  the  (A)  and 
most  effective  part  of  the  oxidizing  flame,  Reducing 

j  r>      r    xi_  a  A.    A  (B)Flames. 

and  B  of  the  reducing  flame.    At  A  metals 

are  oxidized,  and  at  B  oxygen  compounds  are  reduced. 

(See  Part  II,  Exps.  97,  101.) 


ILLUMINATING   GASES  191 

Sometimes  a  tapering  tube  with  a  small  opening  at 
one  end,  called  a  blowpipe,  is  used  to  produce  these  flames. 
A  special  tube  with  a  flattened  top  is  put  inside  the  burner 
tube  to  produce  a  luminous  flame.  The  tip  of  the  blow- 
pipe rests  in  or  near  this  flame,  and  if  air  is  gently  and 
continuously  blown  through  the  blow- 
pipe, a  long,  slender  flame  is  produced, 
called  a  blowpipe  flame  (Fig.  62).  It 
is  like  the  Bunsen  flame  as  far  as  its  Fig.  62.  —  Blowpipe 
oxidizing  and  reducing  properties  are  Flame,  Showing  Ox- 

to,        ™,  .  idizing  (A)  and  Re- 

concerned.     The    mouth   blowpipe    is       ducing  (B)  parts. 
used  in  the  laboratory  and  by  jewelers 
and  mineralogists.     In  the  laboratory  and  in  some  indus- 
trial processes  powerful  blowpipes  are  used.    Air  is  forced 
through  the  apparatus  by  bellows  or  an  air-compression 
machine,  and  the  rapid  combustion  of  the  illuminating 
gas  produces  a  hot,  powerful  flame. 

EXERCISES 

1.  Describe  the  manufacture  of  coal  gas.       Draw  a  diagram  of  the 
apparatus. 

2.  Apply  Exercise  i  to  water  gas. 

3.  Essay  topics:     (a)   The    by-products    of  co?l    gas    manufacture. 
(b)  Manufacture  of  Welsbach  mantles,     (c)  Illumination  in  light  houses. 
(See  National  Geographic  Magazine,  January,  1913.) 

4.  Give  the  equation  for  the  interaction  of  carbon  and  steam. 

6.  What  is  a  flame?  Illustrate  your  answer.  Describe  the  structure  of 
a  candle  flame.  What  are  the  chief  gaseous  products  of  combustion?  Why 
do  lamps  sometimes  smoke? 

6.  Draw  a  diagram  of  the  parts  of  a  candle  flame  from  actual  observation. 

7.  Apply  Exercise  6  to  an  illuminating  gas  flame. 

8.  Sketch  a  Bunsen  flame,  showing  the  oxidizing  and  reducing  parts. 

9.  Describe  the  Welsbach  burner  and   light.     What    is    the  mantle? 
What  is  its  relation  to  the  light? 

10.  Explain:    (a)  "This  is  a  19  candle  power  gas."     (6)  "Water  gas 
is  a  carburetted  gas."    (c)  "Large  oil  lamps  have  a  central  draft." 

11.  Home  exercises:    (a)  Examine  a  Welsbach  burner  from  which  the 


IQ2  CHEMISTRY 

mantle  has  been  removed  and  compare  with  a  Bunsen  burner,  (ft)  Examine 
the  burner  and  flame  of  a  gas  cooking  range.  Compare  both  with  a  Bunsen 
burner  and  flame,  (c)  Examine  a  kerosene  lamp  and  sketch  the  burner. 
(d)  Test  illuminating  gas  for  sulphur  compounds  (274). 

PROBLEMS 

1.  A  candle  weighing  50  gm.  consists  of  a  wax  composed  of  88  per  cent 
carbon  and  .12  per  cent  hydrogen.    What  weight  of  carbon  dioxide  and  of 
water  will  be  formed  by  burning  half  the  candle? 

2.  The  capacity  of  a  gas  holder  is  1,000,000  cu.  ft.    Calculate  the  volume 
of  the  ingredients  (see  table  208)  if  the  holder  were  full  of  (a)  coal  gas,  of 
(b)  water  gas,  and  of  (c)  equal  portions  of  these  gases. 

3.  How  much  dry  illuminating  gas  at  10°  C.  and  530  mm.  will  fill  a  tank 
having  a  capacity  of  800  cu.  m.?     (Specific  gravity  of  iLuminating  gas  is 
0.5  referred  to  air  and  a  liter  of  air  weighs  1.293  gm-) 

4.  A  bottle  inverted  over  water  contains  53.2  cc.  of  illuminating  gas  at 
870  mm.  and  18.5°  C.    What  is  the  volume  of  the  dry  gas  under  standard 
conditions? 

5.  How  many  liters  of  hydrogen  and  of  carbon  monoxide  at  10°  C.  and 
750  mm.  will  be  formed  by  passing  100  gm.  of  steam  over  incandescent 
carbon? 

6.  An  acetylene  gas  plant  consumes  100  cubic  feet  an  hour.    How  much 
calcium  carbide  would  be  used  in  a  month  of  30  days,  if  the  gas  is  burned 
an  average  of  5  hours  a  day? 

7.  The  vapor  density  of  a  compound  of  hydrogen  and  carbon  is  .875. 
Its  composition  is  C  =  85.714  and  H  =  14.286.     Calculate  its  molecular 
formula. 

8.  What  volume  of  air  (containing  21  per  cent  of  oxygen  by  volume)  will 
be  required  for  the  combustion  of  100  tons  of  coal,  assuming  that  the  coal 
is  80  per  cent  pure  carbon  and  burns  to  carbon  dioxide? 


CHAPTER  XVII 
OTHER  CARBON  COMPOUNDS  —  FOOD  AND  NUTRITION 

215.  Introduction.  —  Carbon   forms  a  very  large  num- 
ber of  compounds.     Several  organic  compounds,  as  most 
carbon  compounds  are  called,  have  already  been  discussed 
in  Chapters  XV  and  XVI.      A  few  others  will  be  con- 
sidered' in  the  present  chapter,  especially  in  their  relation 
to  food  and  nutrition. 

216.  Composition  of  Organic  Compounds.  —  Although 
the  number  of  organic  compounds  is  very  large,  they 
contain  only  a  few  elements.     Hydrocarbons,  as  already 
stated,   contain  only  carbon  and  hydrogen.     Vegetable 
substances,  typified  by  starch,  sugar,  and  fruit  acids  and 
flavors,  contain  carbon,  hydrogen,  and  oxygen.     Animal 
substances,  like  albumin,  gelatin,  and  lean  meat  contain 
nitrogen  as  well  as  carbon,  hydrogen,  and  oxygen;    some 
also  contain  sulphur  or  phosphorus.    Fats  contain  carbon, 
hydrogen,  and  oxygen.      (See  Part  II,  Exp.  103.) 

Many  organic  compounds  contain  radicals.  These  radicals  are 
groups  of  atoms  analogous  to  hydroxyl  (OH)  and  ammonium  (NHO, 
and  like  other  radicals  they  exist  only  in  combination.  The  radical 
C2H5  is  called  ethyl.  It  is  present  in  many  organic  compounds,  and 
its  presence  in  ordinary  alcohol  gives  rise  to  the  scientific  name, 
ethyl  alcohol.  Methyl  (CH3)  is  another  important  radical,  and 
phenyl  (C6H5)  is  especially  common  in  the  benzene  series  of  organic 
compounds.  The  names  of  many  radicals  are  derived  from  the 
name  of  the  corresponding  hydrocarbon,  e.g.  methyl  from  methane, 
ethyl  from  ethane. 


194  CHEMISTRY 

217.  Classification  of  Organic  Compounds.  -—  In  this 
chapter  we  shall  study  (i)  Carbohydrates,  (2)  Alcohols, 
(3)  Acids,  (4)  Esters,  (5)  Fats,  Glycerin,  and  Soap, 
(6)  Proteins,  (7)  Formaldehyde,  Acetone,  and  Ether. 


Carbohydrates 

218.  Carbohydrates.  —  The    most    important    carbo- 
hydrates  are   the   sugars,   starch,    and    cellulose.     They 
contain   carbon,   hydrogen,   and  oxygen  —  the  last  two 
elements  in  the  ratio  in  which  they  form  water,  hence 
the  term  carbohydrate. 

219.  Sugars.  — The  popular  term  sugar  means  almost 
any  sweet  substance  found  in  fruits,  nuts,  vegetables,  sap 
of  trees,  etc.,  though  it  is  usually  restricted  to  the  ordi- 
nary white  sugar  obtained  from  sugar   cane   and   sugar 
beet.     Chemically,    there    are    many    different    sugars. 
The  most  important  is  ordinary  sugar  or  cane   sugar, 
which  is  also  called  sucrose  or  saccharose.     Other  sugars 
are  dextrose,  levulose,  lactose,  and  maltose. 

220.  Sucrose,  C^H^On,  is  widely  distributed  in  nature. 
Sugar  cane  contains  about  18  per  cent  and  sugar  beets 
from  12  to  15  per  cent.     Considerable  is  also  found  in  the 
sugar  maple,  sorghum,  sweet  fruits,  many  nuts,  blossoms 
of  flowers,  and  honey.     The  source  of  sucrose  is  sugar 
cane  and  sugar  beet. 

Sucrose  is  a  white,  crystallized  solid;  rock  candy  is  well 
crystallized  sugar.  It  is  very  soluble  in  water,  one  part 
of  water  dissolving  about  three  times  its  weight  of  sugar 
at  ordinary  temperatures.  If  sugar  is  carefully  heated 
to  about  160°  C.,  it  melts,  and  on  cooling  forms  a  glassy 
solid.  As  the  temperature  is  raised,  the  sugar  begins  to 
decompose,  and  at  about  210°  C.  water  is  given  off  and 


OTHER   CARBON   COMPOUNDS  195 

a  light  brown  substance  called  caramel  is  formed,  which 
is  used  to  color  soups  and  gravies.  By  further  heating 
a  black  porous  mass  of  carbon  is  finally  obtained,  often 
called  sugar  charcoal.  (See  Part  II,  Exp.  104.) 

221.  The  manufacture  of  sugar  from  sugar  cane  and  sugar  beets 
involves  two  main  operations:    (i)  In  the  preparation  of  raw  sugar 
from  sugar  cane  the  juice  obtained  by  crushing  the  cane  is  first  boiled 
with  a  weak  calcium  hydroxide  solution  to  neutralize  acids,  remove 
impurities,  and  prevent  fermentation,  next  freed  from  excess  of  lime 
by  carbon  dioxide,  and  finally  filtered  through  bone  black.     The 
purified  juice  is  then  evaporated  in  vacuum  pans  until  the  sugar  be- 
gins to  crystallize  from  the  cooled  liquid.     The  crystals  are  then 
separated  from  the  brown  liquid  by  a  centrifugal  machine.     The 
liquid  is  the  familiar  molasses.     In  the  preparation  of  raw  sugar 
from  sugar  beets  the  washed  beets  are  cut  into  slices  and  soaked  in 
water.     The  sugar  dissolves  in  the  water.     The  solution  is  treated  by 
processes  much  like  those  applied  to  cane  sugar  solutions.     (2)  Raw 
sugar  is  dark  colored,  and  must  be  refined  before  it  is  suitable  for  most 
uses,     (a)  The  raw  sugar  is  first  dissolved  in  water,  air  is  blown  in  to 
agitate  the  heated  solution,  and  lime  and  other  substances  are  added 
to  gather  the  impurities  into  a  scum  or  clot.     The  colored  liquid  is 
next  filtered,  first  through  cloth  bags  and  then  through  animal  char- 
coal,    (b)  The  filtered  sirup  is  evaporated  in  large  vacuum  pans  until 
a  sample  shows  that  the  solution  on  cooling  will  deposit  the  right  size 
crystals.     The  crystals  of  sugar  are  separated  from  the  sirup  by  cen- 
trifugal machines.     The  solution  is  boiled  again  to  obtain  more  crys- 
tals or  sold  as  table  sirup.     The  crystals  are  dried  in  a  heated  tube 
called  a  granulator,  so  that  each  grain  will  be  separate.    Hence  the 
name  granulated  sugar. 

222.  Dextrose  and  Levulose.  —  When  sucrose  is  heated 
with  dilute  acids,  the  two  sugars  dextrose  and  levulose  are 
formed.     The  chemical  change  is  an  example  of  hydroly- 
sis and  may  be  represented  thus :  - 

C12H220n     +     H20     =     C6H1206     +     C6H1206 

Sucrose  Dextrose  Levulose 


196  CHEMISTRY 

The  same  change  is  brought  about  by  a  substance  called 
invertase  (see  enzymes,  below).  Dextrose  is  a  white 
solid  about  three  fifths  as  sweet  as  sucrose.  It  is  very 
soluble  in  water,  but  crystallizes  from  it  with  difficulty. 
Dextrose  is  found  in  honey  and  in  many  fruits,  especially 
grapes,  and  is  sometimes  called  grape  sugar.  Another 
name  for  it  is  glucose.  The  thick  sirup  commercially 
called  glucose  contains  about  40  to  50  per  cent  of  dextrose. 
It  is  manufactured  by  heating  starch  with  dilute  sul- 
phuric acid;  if  the  process  is  carried  far  enough,  the 
product  is  a  hard,  waxlike  solid  known  as  commercial 
grape  sugar,  which  is  almost  pure  dextrose.  Glucose  is 
an  inexpensive  substitute  for  sucrose  and  is  extensively 
used  in  making  candy,  jellies,  sirups,  and  other  sweet 
mixtures.  Levulose  is  also  a  sweet,  white  solid  found  in 
fruits  and  honey,  and  is  often  associated  with  dextrose. 
It  is  sometimes  called  fructose  or  fruit  sugar. 

Dextrose,    and   also   levulose,   is    converted    by   yeast 
into  ethyl  alcohol  and  carbon  dioxide,  thus:  - 

C6Hi2O6  =    2C2H5OH   +        2CO2 

Dextrose          Ethyl  Alcohol         Carbon  Dioxide 

This  chemical  change  is  an  example  of  fermentation,  i.e. 
the  conversion  of  an  organic  compound  into  simpler 
substances  by  the  action  of  minute  living  organisms 
called  ferments,  or  of  the  products  secreted  by  them. 
These  products  are  called  enzymes.  The  alcoholic 
fermentation  of  dextrose  is  due  to  the  enzyme  called 
zymase  (235).  Other  enzymes  produce  other  kinds  of 
fermentation,  e.g.  invertase  transforms  sucrose  into 
dextrose  and  levulose.  Dextrose  and  levulose  are  reduc- 
ing agents.  An  alkaline  solution  of  dextrose  is  used  to 
reduce  a  silver  solution  and  deposit  the  silver  as  a  bright 


OTHER   CARBON   COMPOUNDS  197 

film  in  making  reflectors,  mirrors,  Dewar  flasks,  and 
thermos  bottles.  It  also  reduces  a  strongly  alkaline 
mixture  of  copper  sulphate  and  sodium  potassium  tar- 
trate,  known  as  Fehling's  solution.  When  this  solution 
is  boiled  with  dextrose  (or  a  reducing  sugar),  a  reddish 
copper  compound  (cuprous  oxide,  Cu20)  is  formed.  This 
experiment  is  often  used  as  a  test  for  dextrose  and  similar 
sugars.  Solutions  of  dextrose  and  levulose  rotate  the 
plane  of  polarized  light  —  dextrose  to  the  right  and 
levulose  to  the  left.  That  is,  when  their  solutions  are 
placed  in  a  sugar-polariscope  and  examined,  the  light 
instead  of  passing  entirely  through  the  instrument  is 
extinguished;  and  in  order  to  bring  about  illumination 
again,  the  plane  of  the  polarized  light  must  be  rotated  a 
certain  number  of  degrees  in  order  to  compensate  for  the 
rotation  caused  by  the  sugar  solution.  By  means  of  this 
instrument  valuable  information  can  be  obtained  about 
the  kind  and  proportion  of  sugar  in  solutions.  (See  Part 
II,  Exps.  105,  106,  107,  122.) 

223.  Isomerism.  —  The  formula  of  both  dextrose  and  levulose 
is  C6Hi2O6,  yet  their  properties  are  different.     The  difference  is  due 
to  a  different  arrangement  of  the  atoms  in  a  molecule.     Such  com- 
pounds are  called  isomers  and  illustrate  isomerism.     There  are  many 
cases  of  isomerism  among  organic  compounds,  especially  the  carbo- 
hydrates. 

224.  Lactose   occurs  in  the  milk  of  mammals  and  is 
sometimes   called   milk   sugar.     It   gives  milk  its  sweet 
taste.     Cow's  milk  contains  from  3  to  5  per  cent  of  lactose. 
Crystallized  lactose    (Ci2H220ii.H2O)    is   a   rather   hard, 
gritty  solid,  much  like  sucrose,  though  not  so  sweet  or 
soluble.     A  solution  of  lactose  turns  the  plane  of  polarized 
light  to  the  right  and  reduces  Fehling's  solution.     Lactose 
is  not  fermented  by  ordinary  yeast,  but  a  special  ferment, 


198  CHEMISTRY 

called  lactic  ferment,  converts  it  into  alcohol  and  lactic 
acid.  The  lactic  acid  gives  milk  its  sour  taste  and  also 
assists  in  curdling  the  milk,  i.e.  in  changing  the  casein 
into  a  clot  or  curd.  Lactose  is  obtained  from  whey, 
which  is  the  liquid  left  after  the  solids  have  been  pressed 
from  milk  curdled  by  rennet  in  the  manufacture  of  cheese. 
Lactose  is  used  in  preparing  infant  foods  and  certain 
medicines. 

225.  Maltose  is  formed  from  starch  by  malt,  hence 
the  name  maltose.     The  transformation  is  caused  by  the 
enzyme  diastase.     Malt  is  prepared  by  allowing  moist 
barley  to  sprout  in  a  warm  place;    the  diastase  forms 
during    this    process.     Maltose    ferments    readily    with 
yeast,  forming  alcohol  and  carbon  dioxide,  and  is  manu- 
factured in  large  quantities  for  the  commercial  produc- 
tion of  alcohol  and  fermented  liquors  (235).    With  dilute 
acids,     maltose    forms     dextrose    by    hydrolysis.     Like 
lactose,  maltose  is  a  sweet  solid,  very  soluble  in  water, 
from  which  it  forms  crystals  (C^H^On.H^O);  its  solution 
turns  the  plane  of  polarized  light  to  the  right  and  reduces 
Fehling's  solution. 

226.  Starch   is   a   widely  distributed  and  very  abun- 
dant carbohydrate.     It  is  found  in  wheat,  corn,  and  all 
other  grains;   in  potatoes,  beans,  peas,  and  similar  vege- 
tables;  and  in  large  quantities  in  rice,  sago,  tapioca,  and 
nuts.     Many  parts  of  plants  contain  starch;  for  example, 
the  stalk,  stem,  leaves,  root,  seed,  and  fruit.     The  food 
value  of  vegetables  depends  largely  on  the  starch  they 
contain.     Very  large  quantities  of  starch  are  consumed 
as  food;   much  is  used  in  laundries,  paper  manufactories, 
and  cotton  cloth  mills,  and  in  the  manufacture  of  glucose, 
alcphol  and  alcoholic  beverages,  and  adhesives. 

Starch,  as  usually  seen,  is  a  white  mass,  but  really 


OTHER   CARBON   COMPOUNDS  199 

consists  of  minute  grains  which  vary  with  the  plant,  as 
may  be  seen  by  examining  starch  with  a  microscope 
(Fig.  63).  Starch  is  extracted  from  many  plants  —  in 
the  United  States  largely  from  corn  and  wheat  and  in 
Europe  chiefly  from  potatoes,  rice,  and  wheat.  Starch 


1^   SJW 

^•.-X 

Fig.  63.  — Starch  Grains  (Magnified)  —  Wheat  (Left),  Rice  (Center) 
Corn  (Right). 

is  only  very  slightly  soluble  in  water  because  the  granules 
are  enveloped  in  an  insoluble  membrane  of  cellulose 
(229).  But  if  boiled  with  water,  the  membrane  bursts, 
the  grains  swell,  dissolve  to  some  extent,  and  form  a  jelly- 
like  mass  —  the  familiar  starch  paste.  Starch  gives  a  blue 
colored  substance  when  added  to  iodine  solution,  and  its 
presence  in  many  vegetables  and  foods  can  be  readily 
shown  by  grinding  the  substance  in  a  mortar  with  warm 
water  and  adding  a  drop  of  dilute  iodine  solution.  It 
does  not  ferment  nor  reduce  Fehling's  solution.  Starch 
is  a  complex  carbohydrate  and  its  composition  corre- 
sponds to  the  formula  (C6Hi0O5)x.  Starch  is  readily 
transformed  into  other  carbohydrates.  Thus,  with  cer- 
tain enzymes  it  forms  maltose  and  dextrin  (225,  235), 
while  with  dilute  acids  it  hydrolyzes  into  maltose,  dex- 
trin, and  glucose  (222).  (See  Part  II,  Exps.  108, 123, 124.) 

Wheat  flour  contains  about  70  per  cent  of  starch.  The  remainder 
is  chiefly  water  and  gluten,  though  small  quantities  of  mineral  matter 
and  fat  are  present.  In  making  bread,  the  flour,  water,  and  yeast 
are  thoroughly  mixed  into  dough,  which  is  put  in  a  warm  place  to 


200  CHEMISTRY 

rise.  Fermentation  begins  at  once.  Enzymes  from  the  yeast  change 
the  starch  into  dextrose,  or  a  similar  fermentable  substance,  which 
undergoes  fermentation,  forming  alcohol  and  carbon  dioxide.  The 
gases  escape  in  part  through  the  dough,  which  becomes  light  and 
porous.  When  the  dough  is  baked,  the  heat  kills  the  yeast  plant 
and  fermentation  stops;  but  the  alcohol,  carbon  dioxide,  and  some 
water  escape  and  puff  up  the  mass  still  more.  The  heat,  however, 
soon  hardens  the  starch,  gluten,  etc.,  into  a  firm,  porous  loaf. 

227.  Glycogen,  (CeHioOs)*,  is  a  white,  amorphous  solid  without 
taste  or  odor,  which  resembles  starch  in  some  respects.  It  occurs 
abundantly  in  the  liver  and  in  smaller  proportions  in  muscles,  blood, 
etc.  The  liver  acts  as  a  sort  of  storehouse  for  glycogen,  which  is 
doled  out  to  the  body  as  needed.  Glycogen  plays  somewhat  the  same 
role  in  animals  as  starch  does  in  plants,  and  is  sometimes  called 
"animal  starch."  Glycogen  forms  an  opalescent  solution  with  water, 
is  converted  into  maltose,  dextrose,  and  dextrin  by  certain  enzymes, 
gives  a  red  color  with  iodine  solution,  and  hydrolyzes  to  dextrose 
with  acids;  it  does  not  reduce  Fehling's  solution  nor  ferment  with 
yeast. 


228.  Dextrin  (CsiH^Osi  probably)  is  a  light  brown  or 
white  sweetish  solid  formed  by  heating  starch  to  200°- 
250°  C.     It  dissolves  in  cold  water  and  forms  a  sticky 
solution   which   is   used   as    an   adhesive,    especially   on 
postage  stamps.     It  is  also  used  in  making  mucilage,  for 
thickening  colors  and  sticking  them  to  cloth  in  calico- 
printing,  and  as  an  ingredient  of  candies  and  beverages. 
When  starched  clothes  are  ironed,  the  hot  iron  changes 
some  of  the  starch  into  dextrin,  which  gives  a  gloss  to 
the  fabric.      Dextrin  is  also    formed   from  starch  when 
bread  is  baked  or  toasted.     There  are  several  dextrins. 
(See  Part  II,  Exp.  125.) 

229.  Cellulose,  (C6Hi005)x,  is  the  substance  of  the  cell 
walls  of  which  plants  are  made,  and  is  therefore  very 
widely  distributed.     Wood  contains  cellulose  and  related 
compounds,  while  cotton,  linen,  and  the  best  qualities 


OTHER   CARBON   COMPOUNDS  201 

of  filter  paper  are  nearly  pure  cellulose.  Pure  cellulose 
is  a  white  substance,  insoluble  in  most  liquids,  but  solu- 
ble in  a  mixture  of  ammonia  and  copper  hydroxide.  Con- 
centrated sulphuric  acid  dissolves  it  slowly;  and  if  the 
solution  is  diluted  and  boiled,  the  cellulose  is  changed 
into  a  mixture  of  glucose  and  dextrin.  Sulphuric  acid 
of  a  special  strength,  if  quickly  and  properly  applied 
to  paper,  changes  it  into  a  tougher  form  called  parch- 
ment paper.  The  latter  is  often  substituted  for  animal 
parchment  (e.g.  sheepskin),  and  has  a  variety  of  uses. 

230.  Derivatives  of  Cellulose.  —  With  nitric  acid  cellulose  forms 
cellulose  nitrates.     One  of  the  cellulose  nitrates  is  gun  cotton.     It 
looks  like  ordinary  cotton,  and  may  be  spun,  woven,  and  pressed 
into  cakes.     It  burns  quickly,  if  unconfined,  but  when  ignited  by  a 
percussion  cap  or  when  burned  in  a  confined  space,  gun  cotton  explodes 
violently.     It  is  used  in  blasting  and  for  torpedoes  and  submarine 
mines.     A  mixture  of  gun  cotton,  ether,  and  alcohol  soon  becomes 
a  plastic  mass,  which  upon  being  rolled  and  carefully  dried  forms  a 
transparent  solid;    this  substance  is  called  smokeless  powder,  and 
when  exploded  forms  carbon  dioxide  and  monoxide,  nitrogen,  hydro- 
gen, and  water  vapor  —  all  colorless  gases.     A  solution  of  certain 
cellulose  nitrates  in  a  mixture  of  alcohol  and  ether  is  called  collodion. 
When  collodion  is  poured  or  brushed  upon  a  glass  plate  or  the  skin, 
the  solvent  evaporates,  leaving  behind  a  thin  film.     It  is  used  in 
preparing  photographic  films  and  as  a  coating  for  wounds.     A  mixture 
of  camphor  and  cellulose  nitrates  is  called  celluloid,  which  is  widely 
used  in  making  photographic  films  and  as  a  substitute  for  ivory. 
(See  Part  II,  Exp.  126.) 

231.  Paper  consists  chiefly  of  cellulose  matted  together.     Most 
paper,  especially  that  used  for  newspapers,  is  made  from  wood. 
Considerable  writing  paper,  however,  is  still  made  from  cotton  and 
linen  rags.     In  making  paper  from  wood,  the  latter  is  reduced  to  a 
pulp,  which  is  washed,  spread  on  a  frame  or  an  endless  wire  gauze, 
partly  dried,  and  pressed  by  rollers  into  a  compact  sheet.     The  pulp 
is  prepared  mechanically  by  grinding  the  wood  upon  a  revolving 
stone  or  chemically  by  heating  it  under  pressure  with  sodium  hydrox- 


202  CHEMISTRY 

ide  or  calcium  bisulphite  (acid  calcium  sulphite).     Chemical   pulp 
has  longer  and  stronger  fibers  than  mechanical  pulp. 

232.  Compounds    related    to    Carbohydrates.  —  Fruits    contain 
pectocellulose,  which  is  a  compound  of  cellulose  and  a  complex  car- 
bohydrate called  pectin.     When  certain  fruits,  especially  those  not 
quite  ripe,  are  boiled  with  water,  the  pectin  undergoes  hydrolysis 
and  forms  pectic  acid  and  its  salts,  which  set  to  a  jellylike  mass  on 
cooling.     The  making  of  jellies  from  currants,  grapes,  apples,  and 
other  fruits  depends  on  the  presence  and  transformation  of  pectin. 
Fruits  containing  much  acid  lose  their  jellying  property  if  boiled  too 
long,  owing  to  the  decomposition  of  the  pectic  acid,  while  fruits  which 
are  too  ripe  do  not  form  jelly  because  they  contain  only  a  small 
amount  of  pectin. 

Alcohols 

233.  Methyl  Alcohol,  CH4O  or  better  CH3.OH,  is   a 
colorless  or  slightly  yellowish  liquid,  much  like  ordinary 
alcohol.    It  boils  at  about  66°  C.,  and  burns  with  a  pale 
blue  flame  which  deposits  no  soot.     It  mixes  with  water 
in  all  proportions.     It  is  cheaper  than  ethyl  alcohol,  and 
is  used  as  a  fuel,  a  solvent  for  fats,  oils,  and  shellac,  and 
in  the  manufacture  of  varnishes,   dyestuffs,   and  many 
chemicals.     Methyl  alcohol  is  often  called  wood  alcohol 
or  wood  spirit,  because  it  is  one  of  the  products  obtained 
by  the   dry  distillation   of  wood   (176).      It  has  a  dis- 
agreeable   odor    and    is    poisonous.     The    concentrated 
liquid  causes  blindness  and  even  death. 

234.  Ethyl    Alcohol,    C2H6O    or    better    C2H5.OH,    is 
ordinary   alcohol,   and  is   often   called   grain   alcohol   or 
simply  alcohol.     It  is  a  colorless,  volatile  liquid,  with  a 
pleasant    odor.     It    is    lighter    than    water,    its    specific 
gravity  being  about  0.8.     It  boils  at  about  78°  C.  and 
solidifies  at  about  —  112°  C.    It  burns  with  a  hot,  nearly 
colorless,  non-smoking  flame,  and  is  sometimes  used  as  a 
source  of  heat.     Alcohol  mixes  with  water  in  all  propor- 


OTHER   CARBON   COMPOUNDS  203 

tions.  The  ordinary  commercial  variety  contains  about 
95  per  cent  of  alcohol.  Pure  or  absolute  alcohol  is  ob- 
tained by  removing  the  remaining  water  with  lime  and 
dehydrated  copper  sulphate.  Small  quantities  of  ordinary 
alcohol  produce  intoxication,  while  large  quantities  are 
poisonous.  Alcohol  is  an  excellent  solvent  for  gums, 
oils,  and  resins,  and  is  extensively  used  in  the  manufac- 
ture of  varnishes,  essences,  extracts,  tinctures,  perfumes, 
and  medicines.  Many  organic  compounds,  as  ether  and 
chloroform,  are  prepared  from  alcohol.  Some  vinegar 
is  made  from  alcohol.  In  museums  alcohol  is  used  to 
preserve  specimens.  Denatured  alcohol  is  essentially  a 
mixture  of  100  parts  ethyl  alcohol,  10  parts  methyl  al- 
cohol, and  a  small  proportion  of  benzine,  pyridine  (or  a 
similar  mixture);  it  is  not  taxed,  like  ethyl  alcohol,  and 
in  its  legalized  forms  is  used  as  a  cheap  substitute  for 
ordinary  alcohol.  It  is  unfit  for  drinking,  largely  on 
account  of  the  disagreeable  taste,  but  is  suitable  for  in- 
dustrial uses.  (See  Part  II,  Exp.  109.) 

235.  Manufacture  of  Alcohol.  —  Alcohol  is  produced  by  the 
fermentation  of  certain  sugars.  Alcoholic  fermentation  is  caused  by 
the  enzyme  zymase  that  is  secreted  by  ordinary  yeast.  When  yeast 
is  added  to  a  solution  of  dextrose,  maltose,  or  any  other  fermentable 
sugar,  the  yeast  plants  multiply  rapidly.  The  chemical  changes 
are  numerous  and  complex,  but  the  main  products  resulting  from 
the  action  of  the  enzyme  from  the  yeast  upon  dextrose  and  maltose 
are  alcohol  and  carbon  dioxide,  thus:  — 

C6H1206     =     2C2H60     +       2CO2 

Dextrosd  Alcohol  Carbon  Dioxide 

Ci2H22Ou  +  H2O  =  4C2H6O  +  4CO2 

Maltose  Alcohol 

Alcohol  is  made  chiefly  from  the  starch  obtained  from  corn  and 
potatoes.  Malt  is  added  to  a  warm  mixture  of  starch  and  water, 
and  the  enzyme  diastase  in  the  malt  converts  the  starch  into  maltose. 


204  CHEMISTRY 

When  the  action  is  over,  this  mixture  is  cooled  and  diluted,  and 
yeast  is  added.  The  zymase  in  the  yeast  changes  the  maltose  into 
alcohol  and  carbon  dioxide.  The  alcohol,  which  constitutes  about 
15  per  cent  of  the  final  mixture,  is  separated  by  distillation.  (See 
Part  II,  Exp.  127.) 

236.  Alcoholic   Beverages.  —  Wine,   beer,    and    distilled   liquors 
are  essentially  mixtures  of  alcohol  and  water  made  by  the  fermenta- 
tion of  sugars.     They  differ  mainly  in  the  proportion  of  alcohol.     The 
particular  flavor  is  due  to  small  quantities  of  different  substances 
which  are  intentionally  added,  obtained  from  the  raw  materials,  or 
formed  by  special  processes  of  manufacture.     Beer  contains  from 
3  to  7  per  cent  of  alcohol,  wines  from  6  to  20,  rum,  brandy,  and  whisky 
from  40  to  60  or  more  per  cent. 

Acids 

237.  Acetic  Acid,   C2H4O2   or    CH3.COOH.  —  This    is 
the  most  common  organic  acid.     It  is  manufactured  on 
a  large  scale  by  the  dry  distillation  of  wood  (176).     The 
dark-red  watery  distillate,  which  is  called  pyroligneous 
acid,  contains  about  10  per  cent  of  acetic  acid,  besides 
methyl  alcohol  and  acetone   (253).     Very  concentrated 
acetic  acid  is  called  glacial  acetic  acid,  because  at  about 
17°  C.  it  becomes  an  icelike  solid. 

Commercial  acetic  acid  is  a  water  solution  containing 
about  30  per  cent  of  pure  acetic  acid.  It  is  a  colorless 
liquid,  having  a  rather  pungent  odor.  It  is  a  weak  acid. 
Acetic  acid  is  used  to  prepare  acetates,  dyestuffs,  medi- 
cines, and  white  lead.  Some  of  its  salts  —  the  acetates 
—are  useful  compounds,  e.g.  lead  acetate  and  Paris  green 
(see  these  salts).  (See  Part  II,  Exps.  110,  111,  112.) 

238.  Vinegar  is  dilute  acetic  acid,  containing  from  4  to  6  per  cent 
of  the  acid.     It  is  prepared  by  oxidizing  dilute  alcohol,  the  essential 
change  being  represented  thus:  — 

C2H60     +     02     =     C2H4O2     +     H2O 

Alcohol  Oxygen  Acetic  Acid  Water 


OTHER   CARBON   COMPOUNDS 


205 


The  transformation  is  accomplished  by  fermentation.  When  dilute 
solutions  of  alcohol,  e.g.  beer  or  weak  wines,  are  exposed  to  air,  they 
slowly  become  sour,  owing  to  the  conversion  of  alcohol  into  acetic 
acid.  The  change  is  caused  by  the  presence  and  activity  of  a  ferment 
known  as  mycoderma  aceti,  or  "mother  of  vinegar."  Strong  wines 
and  pure  dilute  alcohol  do  not  become  sour,  because  the  ferment  can- 
not live  in  such  liquids.  Substances  containing  starch  or  ferment- 
able sugars,  e.g.  fruit  juices,  cider,  and  molasses,  slowly  ferment 
when  exposed  to  the  air  (which  always  contains  the  organisms  neces- 
sary for  the  chemical  transformations),  forming  alcohol  first  and 
finally  vinegar.  Cider  vinegar  is  made  this  way.  Vinegar  is  often 
made  by  the  quick  vinegar 
process.  Dilute  impure  alco- 
hol is  introduced  at  the  top  of 
a  vat  filled  with  beechwood 
shavings  soaked  in  old  vine- 
gar, trickles  through  the  shav- 
ings, and  collects  at  the  bottom; 
holes  at  the  bottom  and  top 
allow  air  to  enter  and  escape 
freely  (Fig.  64) .  In  its  passage 
it  comes  in  contact  with  the 
ferment  and  oxygen,  and  is 
partially  converted  into  vine- 
gar. The  operation  is  re- 
peated until  the  change  is  complete. 


Fig.  64. 


Apparatus    for  Manufactur- 
ing Vinegar. 


Thus  prepared,  the  vinegar 
lacks  the  flavor,  odor,  and  color  of  cider  vinegar,  but  these  deficiences 
may  be  artificially  supplied.  Vinegar  is  used  as  a  condiment  and  in 
making  pickles  and  similar  relishes. 

239.  Butyric  Acid,  C4H8O2,  is  the  acid  which  gives  the 
disagreeable  odor  to  rancid  butter.  A  derivative  occurs 
in  sweet  butter.  Simple  derivatives  of  stearic  acid 
(Ci8H36O2)  and  palmitic  acid  (Ci6H32O2)  are  found  in 
beef  suet,  mutton  fat,  butter,  and  other  fats;  these  two 
acids  are  white  solids.  Oleic  acid  (Ci8H34O2)  is  an  oily 
liquid;  certain  derivatives  are  found  in  olive  oil  and 
many  other  oils  and  fats. 


206  CHEMISTRY 

240.  Oxalic  acid  occurs  as  a  salt  in  rhubarb  and  sorrel.     It  is 
very  poisonous.     The  acid  and  some  of  its  salts  decompose  iron  rust 
and  inks  containing  iron,  and  are  often  used  to  remove  such  stains 
from  cloth. 

241.  Lactic  acid,  C3H6O3,  occurs  in  sour  milk,  being  a  product  of 
the  action  of  certain  bacteria  (which  are  in  the  air)  on  the  milk  sugar 
(224).     When  sour  milk  is  used  in  cooking,  the  baking  soda  (HNaCO3) 
and  lactic  acid  interact,  producing  soluble  sodium  lactate  and  carbon 
dioxide  gas. 

242.  Malic  acid,  C4H6O5,  is  found  free  or  as  salts  in  apples,  pears, 
cherries,  currants,  gooseberries,  rhubarb,  grapes,  and  berries  of  the 
mountain  ash  tree;  also  in  the  roots,  leaves,  and  seeds  of  many  vege- 
tables. 

243.  Tartaric  acid,  C4H6O6,  occurs  as  the    potassium 
salt  (acid  potassium  tartrate,  HKC4H4O6)  in  grapes  and 
other  fruits.     During   the   fermentation   of  grape  juice, 
the  impure  salt  is  deposited  in  the  wine  casks.     From 
this  argol  or  crude  tartar  the  acid  is  prepared.     Tar- 
taric acid  is  a  white  crystallized  solid,  soluble  in  water 
and   alcohol.     It  is  used  in  dyeing,  and  as  one  ingredi- 
ent of  Seidlitz  powders.     In  these  and  similar  mixtures 
it   serves   to  decompose   the   other  ingredient,   which  is 
a  carbonate  (see  Sodium  Bicarbonate). 

Purified  acid  potassium  tartrate  obtained  from  argol 
is  commonly  known  as  cream  of  tartar.  It  is  extensively 
used  in  the  manufacture  of  tartrate  baking  powders. 
These  are  mixtures  of  cream  of  tartar  and  sodium  bicar- 
bonate (HNaCO3),  and  starch  (369,  437).  The  starch 
is  added  mainly  to  protect  the  other  ingredients  from 
moisture  and  thus  prevent  the  baking  powder  from 
deteriorating.  When  dissolved  in  water  or  moistened 
by  the  water  in  a  food  mixture,  the  two  ingredients 
interact  and  liberate  carbon  dioxide  as  the  main 
product,  thus:  — 


OTHER   CARBON   COMPOUNDS  207 

HKC4H406  +  HNaC03  =  CO2  +    NaKC4H4O6  +  H2O 

Acid  Potassium  Sodium  Carbon        Sodium  Potassium       Water 

Tartrate  Bicarbonate       Dioxide  Tartrate 

The  gas,  if  generated  within  a  soft  mixture,  escapes 
slowly  and  puffs  up  the  mass,  which  is  ultimately  baked 
to  a  more  or  less  porous  loaf.  (See  Part  II,  Exp.  113.) 

244.  Citric  acid,  C6H8O7,  occurs  abundantly  in  lemons  and  oranges, 
and  in  small  quantities  in  currants,  gooseberries,  raspberries,  and 
other  acid  fruits.    The  acid  is  used  in  making  lemonade. 

Esters 

245.  Esters  are  compounds  of  carbon,  hydrogen,  and 
oxygen   closely   related    to   alcohols   and   organic   acids. 
Thus,  when  ethyl  alcohol  and  acetic  acid  are  mixed  with 
concentrated  sulphuric  acid  and  warmed,  ethyl  acetate 
is  formed.     The  equation  for  the  reaction  is:  - 

C2H5.OH   +  CH3.COOH  =  CH3.COOC2H5  +  H2O 

Ethyl  Alcohol          Acetic  Acid  Ethyl  Acetate  Water 

Ethyl  acetate  has  a  pleasant,  fruitlike  odor,  and  its  forma- 
tion in  this  way  is  a  simple  test  for  alcohol  or  acetic  acid. 
Ethyl  acetate  is  analogous  to  sodium  acetate,  i.e.  the 
organic  salt  contains  the  radical  ethyl,  while  the  metallic 
salt  contains  sodium.  Organic  acids  form  many  impor- 
tant esters.  Some  occur  naturally  in  fruits  and  flowers, 
and  in  many  cases  give  the  fragrance  and  flavor.  Others 
are  prepared  artificially  in  large  quantities  and  used  as 
the  characteristic  ingredient  of  flavoring  extracts,  per- 
fumery, and  beverages.  Ethyl  butyrate  has  the  taste 
and  fragrance  of  pineapples,  amyl  acetate  of  bananas, 
amyl  valerate  of  apples,  methyl  salicylate  of  wintergreen. 
(See  Part  II,  Exp.  111.) 


208  CHEMISTRY 

Fats  and  Glycerin  —  Soap 

246.  General  Relations.  —  Natural  fats  and  oils  are 
essentially  mixtures  of  the  esters,  stearin,  palmitin,  and 
olein.  Stearin  and  palmitin  are  solids  at  ordinary  tem- 
peratures, but  olein  is  a  liquid.  Hence  hard  fats  are 
largely  stearin  and  palmitin,  while  soft  or  liquid  fats  and 
oils  are  largely  olein.  These  three  compounds  are  esters 
derived  from  their  corresponding  acids  by  replacing 
hydrogen  of  the  acid  by  the  radical  of  the  alcohol  glycerin. 
The  radical  of  glycerin  is  glyceryl,  C3H5.  Hence  stearin 
is  glyceryl  stearate,  palmitin  is  glyceryl  palmitate,  and 
olein  is  glyceryl  oleate.  When  fats  are  heated  with  very 
hot  steam  or  with  sulphuric  acid,  they  are  changed  into 
glycerin  and  the  corresponding  acids.  Thus,  with  stearin 
the  change  is  - 


=  C3H5(OH)3  +  3Ci7H35.COOH 

Stearin  Glycerin  Stearic  Acid 

But  if  fats  are  boiled  with  sodium  hydroxide  or  a  similar 
alkali,  glycerin  and  an  alkaline  salt  of  the  corresponding 
acid  are  formed.  Soap  is  a  mixture  of  such  alkaline  salts. 
In  a  few  words,  the  general  relations  are  these:  (i)  fats 
are  esters;  (2)  treated  with  steam  or  acid,  fats  form 
glycerin  and  organic  acids;  (3)  treated  with  alkalies,  fats 
form  glycerin  and  soap. 

247.  Natural  Fats  and  Oils.  —  These  are  often  complex  mixtures. 
The  chief  ingredients,  as  already  stated,  are  stearin,  palmitin,  and 
olein;  small  quantities  of  similar  esters  are  usually  present  and  often 
give  certain  fats  characteristic  properties.  Tallow  is  about  66  per 
cent  stearin  and  palmitin  and  33  per  cent  olein;  it  is  a  solid  fat 
obtained  from  the  sheep  and  ox  and  is  used  in  making  soap  and  candles. 
Lard  consists  of  olein,  stearin,  palmitin,  and  a  little  linolein;  it  is 
obtained  from  the  fat  of  the  hog  and  is  used  in  cooking.  Olive  oil, 
which  is  obtained  from  the  fruit  of  the  olive  tree,  contains  about  72 


OTHER   CARBON   COMPOUNDS  209 

per  cent  of  olein  (and  a  similar  fat)  and  about  28  per  cent  of  stearin 
and  palmitin,  together  with  small  quantities  of  linolein  and  other 
substances;  the  best  qualities  are  used  as  salad  oil,  while  cheap  kinds 
are  utilized  in  cooking,  as  a  lubricant,  and  for  making  soap.  Cotton- 
seed oil  is  similar  to  olive  oil;  it  is  obtained  from  the  seeds  of  the 
cotton  plant  and  is  used  extensively  in  cooking,  as  salad  oil,  and  as 
a  substitute  for  other  oils.  Butter  fat  is  about  60  per  cent  olein,  30 
per  cent  stearin  and  palmitin,  and  5  per  cent  butyrin  (glyceryl  buty- 
rate),  together  with  small  quantities  of  esters  corresponding  to  capric, 
caprylic,  and  myristic  acids.  The  pleasant  flavor  of  butter  is  mainly 
due  to  the  esters  and  some  other  substances  that  are  present  in  small 
proportions.  The  total  amount  of  fat  in  butter  varies  from  about 
78  to  94  per  cent;  according  to  the  Federal  Pure  Food  Law  butter 
offered  for  sale  must  contain  not  less  than  82.5  per  cent  of  fat.  Be- 
sides fat,  butter  contains  water,  casein,  lactose,  and  salt.  Butter 
is  made  by  churning  cream  from  cow's  milk;  this  operation  causes 
the  fat  globules  to  coalesce.  Most  butter  is  colored  artificially, 
tumeric,  annatto,  and  carrot  juice  being  used  for  this  purpose.  Ran- 
cid butter  is  produced  by  the  transformation  of  fat  (by  the  action  of 
bacteria)  into  butyric  acid  and  other  acids  which  have  a  disagreeable 
taste  and  odor.  Oleomargarine  and  other  substitutes  for  butter 
resemble  real  butter.  They  are  made  from  different  fats  and  oils; 
a  little  butter  is  added  or  the  mixture  is  churned  with  milk  to  impart 
the  desired  flavor.  (See  Part  II,  Exp.  114.) 

248.  Glycerin,  C3H8O3  or  C3H5.(OH)3,  is  a  thick,  sweet, 
colorless  liquid.  It  mixes  with  water  and  with  alcohol 
in  all  proportions.  It  absorbs  moisture  readily  from  the 
air.  Heated  in  air  it  decomposes  and  gives  off  irritating 
gases,  like  those  produced  by  burning  fat. 

Glycerin  is  used  to  make  nitroglycerin  (see  below), 
toilet  soaps,  and  printers'  ink  rolls;  it  is  also  used  as 
a  solvent,  a  lubricator,  a  preservative  for  tobacco  and 
certain  foods,  a  sweetening  substance  in  certain  liquors, 
preserves,  and  candy;  as  a  cosmetic;  and  owing  to  its 
non-volatile  and  non-drying  properties,  it  is  used  as  an 
ingredient  of  inks  for  rubber  stamps. 


210  CHEMISTRY 

Glycerin  is  a  by-product  in  the  manufacture  of  soap  and  candles, 
or  it  is  made  directly  by  decomposing  fats  with  steam  under  pressure 
(249). 

As  already  stated,  glycerin  is  an  alcohol,  and  for  this  reason  it  is 
often  called  glycerol.  When  treated  with  a  mixture  of  concentrated 
nitric  and  sulphuric  acids,  it  forms  an  ester  commonly  known  as 
nitroglycerin  (C3H5(ONO2)3).  This  is  a  slightly  yellow,  heavy,  oily 
liquid.  It  is  the  well-known  explosive.  When  kindled  by  a  flame, 
it  burns  without  explosion;  but  if  subjected  to  a  shock  or  heated 
suddenly  by  a  percussion  cap,  it  explodes  violently.  Nitroglycerin 
is  used  in  blasting;  but  since  it  is  dangerous  to  handle  and  transport, 
it  is  usually  mixed  with  some  porous  substance,  such  as  infusorial 
earth,  fine  sand,  clay,  or  even  sawdust.  In  this  form,  it  is  called 
dynamite.  Other  explosives  contain  nitroglycerin,  e.g.  blasting  gela- 
tin and  cordite. 

249.  Soap,  as  already  stated,  is  a  mixture  of  alkaline 
salts  of  organic  acids,  mainly  stearic  and  palmitic  acids. 
Soap  is  made  by  boiling  fats  with  sodium  hydroxide 
or  potassium  hydroxide  solution.  This  process  is  called 
saponification.  Sodium  hydroxide  produces  hard  soap, 
consisting  chiefly  of  sodium  palmitate,  sodium  stearate, 
and  sodium  oleate.  Potassium  hydroxide  produces  a 
soft,  semi-fluid  soap,  which  contains  mainly  the  cor- 
responding potassium  salts.  In  the  case  of  stearin 
(glyceryl  stearate)  the  change  may  be  represented 
thus:- 

C3H5(Ci7H35.COO)3  +  3NaOH  =  3Ci7H35COONa  +  C3H5(OH)3 

Stearin  Sodium  Sodium  Glycerin 

Hydroxide  Stearate 

The  fats  used  in  soap-making  vary.  Tallow,  lard,  palm 
oil,  and  cocoanut  oil  make  white  soaps.  Bone  grease 
or  house  grease,  together  with  tallow,  palm  oil,  cotton- 
seed oil,  and  rosin,  make  yellow  soaps.  Olive  oil  is  used 
for  making  castile  soap.  Sodium  and  potassium  soaps 
are  soluble  in  water.  If  water  contains  dissolved  calcium 


OTHER   CARBON   COMPOUNDS  211 

or  magnesium  salts,  these  interact  with  the  soap  and  form 
insoluble  calcium  or  magnesium  palmitate,  etc.  (423). 

Most  soaps  are  made  by  boiling  the  fat  and  alkali  in  a  huge  kettle. 
This  operation  produces  a  thick,  frothy  mixture  of  soap,  glycerin, 
and  alkali.  At  the  proper  time  salt  is  added,  thereby  causing  the 
soap  to  separate  and  rise  to  the  top.  The  liquid  beneath  is  drawn  off, 
and  from  it  glycerin  is  extracted.  Some  soaps  are  boiled  again  with 
rosin  or  cocoanut  oil,  and  then  mixed  (if  desired)  with  perfume,  color- 
ing matter,  or  filling  material  (such  as  sodium  silicate,  sand,  or  borax). 
Floating  soaps  are  made  by  forcing  air  into  the  semi-solid  mass  before 
cooling.  The  best  quality  soaps  are  prepared  carefully  so  that  the 
finished  product  will  not  contain  any  unchanged  fat  or  "free  alkali," 
i.e.  an  excess  of  sodium  hydroxide.  (See  Part  II,  Exps.  116,  116.) 

The  cleansing  action  of  soap  is  ascribed  to  two  causes,  (i)  Soap 
hydrolyzes  with  water  —  especially  hot  water  —  and  the  liberated 
alkali  (sodium  hydroxide)  acts  upon  the  grease  and  oil  that  is  usually 
mixed  with  the  dirt.  (2)  Soap  causes  oils  to  form  minute  drops,  which 
remain  suspended  in  water  and  can  be  readily  removed.  The  second 
cause  is  the  more  efficient. 

Proteins 

250.  General  Characteristics.  —  Proteins  form  the 
chief  part  of  the  solid  (except  the  mineral  part  of  bone) 
and  liquid  substances  in  the  body,  being  abundant  in 
blood,  tissues,  muscles,  and  nerves;  they  also  occur  in 
most  parts  of  plants,  especially  in  the  seeds.  The  food 
of  all  animals  must  contain  protein  in  some  form  (258, 
259).  The  body  of  the  average  man  is  about  18  per 
cent  protein.  Life  processes  involve  transformations  of 
proteins.  The  term  protein  (formerly  proteid)  includes 
many  very  complex  substances  which  resemble  one 
another  in  composition  and  in  properties.  All  proteins 
contain  nitrogen,  carbon,  oxygen,  and  hydrogen;  most 
also  contain  sulphur,  several  phosphorus,  and  a  few  iron. 
Common  proteins  contain  from  15  to  18  per  cent  of 


212  CHEMISTRY 

nitrogen.  When  burned,  they  produce  a  disagreeable 
odor  and  liberate  ammonia  —  facts  which  may  easily  be 
verified  by  burning  white  of  egg  or  gelatin.  All  proteins 
putrefy,  that  is,  they  decompose  and  evolve  foul  gases, 
such  as  hydrogen  sulphide  and  derivatives  of  ammonia. 
(See  Part  II,  Exps.  117,  118.) 

251.  Groups.  —  Albumins  occur  in  the  white  of  eggs, 
milk,  muscle,  blood,  and  seeds  of  plants,  especially  in 
beans,  peas,  and  lentils,  and  to  some  extent  in  wheat, 
rye,  and  barley.  Albumins  are  soluble  in  pure  cold  water 
and  are  coagulated  by  heat;  the  latter  property  is  readily 
seen  when  an  egg  is  cooked.  They  are  important  in- 
gredients of  food.  Globulins  are  also  found  in  eggs, 
milk,  muscle,  blood,  and  seeds.  The  fluid  part  of  the 
blood  contains  about  0.4  per  cent  of  a  globulin  called 
fibrinogen,  and  from  muscle  are  obtained  myosin  and 
myogen;  the  seeds  of  some  legumes  contain  relatively 
large  per  cents  of  globulins,  e.g.  kidney  beans  20,  peas 
10,  lentils  13.  Globulins  are  relatively  insoluble  in  water 
and  dilute  acids,  but  they  dissolve  in  dilute  solutions  of 
neutral  salts  (e.g.  sodium  chloride).  Glutelins  and  pro- 
tamins,  as  well  as  albumins  and  globulins,  are  ingredients 
of  cereals.  For  example,  gluten,  the  sticky  substance 
in  wheat  flour  that  makes  dough  tenacious  and  elastic, 
contains  glutenin  and  gliadin.  Glutelins  are  insoluble 
in  water  and  neutral  salt  solutions,  but  dissolve  in 
very  dilute  acids  and  alkalies.  Albuminoids  are  related 
to  albumins.  They  make  up  the  framework  of  animal 
tissues  and  are  constituents  of  the  epidermis.  Thus, 
collagen  is  found  in  connective  tissue  and  cartilage,  and 
yields  gelatin  by  boiling  with  water;  the  keratins  are  the 
chief  constituents  of  the  hard  parts  of  the  skin  and  its 
appendages  (e.g.  hair,  hoofs,  nails,  feathers,  etc.),  and 


OTHER   CARBON   COMPOUNDS  213 

elastin  makes  up  the  yellow  elastic  fibers  of  ligaments 
and  the  outer  walls  of  arteries.  Albuminoids  are  among 
the  least  soluble  of  the  proteins.  Phosphoproteins  occur 
especially  in  milk  and  eggs.  Casein  is  the  chief  nitro- 
genous substance  in  milk  and  imparts  the  bluish  color 
to  skimmed  milk.  The  casein  can  be*  separated  from 
milk  as  a  more  or  less  firm  curd  by  dilute  acids.  Thus, 
when  milk  sours,  the  lactic  acid  formed  by  the  fermenta- 
tion of  the  lactose  causes  the  casein  to  separate  as  a  firm 
curd.  Similar  changes  are  produced  by  dilute  sulphuric 
acid  and  rennet.  Rennet  and  certain  preparations 
(e.g.  junket  tablets)  contain  an  enzyme  (rennin)  which 
causes  the  casein  and  fat  to  separate  as  a  firm  mass  from 
which  cheese  can  be  made.  Vitellin  is  an  ingredient  of 
hen's  eggs.  The  most  important  of  the  hemoglobins  is 
ordinary  hemoglobin  of  the  blood.  Hemoglobin  is  a 
compound  of  a  protein  called  globin  and  a  substance 
containing  iron,  called  hematin  (^H^^FeOs  probably). 
Hemoglobin  forms  a  readily  decomposable  compound 
with  oxygen  called  oxyhemoglobin ;  these  are  the  sub- 
stances that  distribute  oxygen  to  the  tissues  of  the  body. 
Hemoglobin  is  a  dark  red  solid  which  forms  about  90 
per  cent  of  the  solid  matter  of  the  red  corpuscles  of  the 
blood.  When  oxygen  of  the  inhaled  air  combines  with 
the  hemoglobin  of  the  blood  in  the  lungs,  the  bright  red 
oxyhemoglobin  is  formed  and  this  oxygenized  blood,  so 
to  speak,  in  circulating  through  the  body  gives  up  its 
oxygen  readily  to  the  tissues;  during  the  circulation, 
carbon  dioxide  —  one  of  the  waste  products  of  vital 
processes  —  is  taken  up  by  the  blood  and  carried  to  the 
lungs  where  it  is  exhaled  along  with  other  gases. 


214  CHEMISTRY 


Formaldehyde  —  Acetone  —  Ether 


252.  Formaldehyde,  CH2O,  is  a  gas,  which  dissolves  in  water. 
It  has  an  irritating  odor.     The  commercial  solution  sold  as  formalin 
contains  40  per  cent  of  formaldehyde.     It  hardens  tissue  and  is  used 
as  a  preservative  in  museums.     It  is  a  convenient  and  efficient  disin- 
fectant.    When  so  used,  the  solution  is  vaporized  in  a  special  appara- 
tus and  the  vapors  conducted  into  the  infected  room.     (See  Part  II, 
Exp.  128.) 

253.  Acetone,  C3H6O,  is  a  colorless  liquid  which  has  an  ethereal 
odor.     It  boils  at  about  56°  C.  and  mixes  in  all  proportions  with  water, 
alcohol,  and  ether.     It  is  used  as  a  solvent  for  fats,  oils,  and  waxes, 
and  in  the  preparation  of  smokeless  powders  and  certain  organic 
compounds.     Acetone  is  one  of  the  products  obtained  by  the  dry 
distillation  of  wood  (176). 

254.  Ethyl  ether,  or  ordinary  ether,  C4Hi0O,  is  a  colorless,  volatile 
liquid,  with  a  peculiar,  pleasing  taste  and  odor.     It  boils  at  35°  C., 
and  the  vapor  is  very  inflammable.     The  liquid  should  never  be 
brought  near  a  flame.     It  is  somewhat  soluble  in  water,  and  it  also 
dissolves  water  to  a  slight  extent.     It  mixes  with  alcohol  in  all  pro- 
portions.    It  is  a  good  solvent  for  waxes,  fats,  oils,  and  other  organic 
compounds.     Its  chief  use  is  as  an  anaesthetic.     Ether  is  manufac- 
tured by  heating  a  mixture  of  ethyl  alcohol  and  sulphuric  acid  in  the 
proper  proportions.     (See  Part  II,  Exp.  129.) 

255.  Miscellaneous.  —  Other   aldehydes   are   benzaldehyde    (oil 
of  bitter  almonds,  C7H6O)  and  vanillin  (C8H803);   both  are  used  as 
flavors.     Alkaloids   are   complex   compounds   found  in  plants;    all 
contain    nitrogen.     They    produce    marked    physiological    effects. 
Theine  or  caffeine  occurs  in  tea  and  coffee,  and  theobromine  in  cocoa 
and  chocolate. 

Food  and  Nutrition 

256.  Food  and  Nutrition.  —  Foods  are  the  substances 
that  supply  the  body  (i)  with  materials  for  its  growth, 
maintenance,    and   repair,    and    (2)    with   energy   for   its 
activities.     In  a  word,  food  supplies  the  body  with  build- 
ing material  and  fuel.     Nutrition  is  the  process,  or  group 


FOOD  AND  NUTRITION 


215 


of  processes,  by  which  tissue  building  and  energy  pro- 
duction are  accomplished. 

257.  Nutrients.  —  The  parts  of  food  that  nourish  the 
body  are  called  nutrients  or  foodstuffs.     Nutrients  are 
derived  chiefly  from  three  groups  of  organic  substances, 
viz.   proteins,   carbohydrates,   and   fats.     The   inorganic 
substances   water   and   mineral   matter   are   also   vitally 
connected  with  nutrition,  and  are  obtained  from  food  or 
directly  from  drinking-water. 

258.  Composition  of  Foods.  —  The  average  composition 
(in  per  cent)  of  the  edible  portion  of  some  common  foods 
is  shown  in  the  Table  of  Composition  of  Foods  (p.  216). 

The  composition  of  white  bread,  butter,  and  milk  is 
strikingly  shown  in  Fig.  65.  (See  Part  II,  Exps.  119,  120, 
122,  123.) 


:VWnier35.3 
I*              Waterll.O     Protein  1.0    F.tSS.O    Ash  3.0 
'Vr^\^.^//                                         /A 

Water  87.0 

.Protein  3.4 

Carho- 

.Cnrbo- 
lydrate* 
53.1 

L 



WH 

TE  BREAD 

BUTTER 

MILK 

nel  ralne                          ^  Fuel  v.  ue 
tlO  Calorie*                       RS  310  Calories 
er  pound                              ^J  per  pound 

^  Fuel  ralue 
SSSSSsSJ  121S  Calorie* 
£££Xl  per  pound 

•  '  '    F 

\§§i§»$i  s 

;  ,  ^  p 

Fig.  65.  —  Composition  of  Bread,  Butter,  and  Milk. 

259.  Nutrition.  —  Most  of  the  nutrients  in  food  must 
be  more  or  less  changed  in  order  to  become  of  use  in 
nutrition.  These  changes  take  place  chiefly  in  the 
digestive  organs  and  constitute  the  process  called  diges- 
tion. After  digestion,  the  modified  foodstuffs  are  absorbed 
and  assimilated,  and  as  the  result  of  extremely  compli- 
cated changes  they  become  a  part  of  the  fluids  and  organs 
of  the  body.  The  long  series  of  complex  changes  that 
take  place  after  absorption  from  the  digestive  organs, 


2l6 


CHEMISTRY 


i.e.    the    building    up   and   tearing  down    processes,  are 
usually  included  by  the  term  metabolism. 


TABLE  OF  COMPOSITION  OF  FOODS' 


Foods 

Water 

Protein 

Fat 

Carbo- 
hydrate 

Mineral 
Matter 

Apples    . 

84.6 

0.4 

0.5 

14.2 

0.3 

Bananas  
Beans  (dried) 

75-3 

12.6 

!-3 

22.  Z 

0.6 
1.8 

22.  0 

^0-6 

0.8 

7.    C 

Beefsteak  (sirloin)    
Celery 

61.9 

QA..Z 

18.6 
i.i 

18.5 

O.I 

2.2 

I.O 
I.O 

Cheese  (cream)   
Codfish  (fresh) 

34-2 
82X 

25-9 
16.3 

33-7 
0.2, 

2.4 

3.8 

O.Q 

Corn  (green)    

EgrgS 

75-4 
72.7 

3-i 
14.8 

I.I 

IOX 

19.7 

0.7 

I.O 

Grapes    

77-4 

1.3 

1.6 

19.2 

0.5 

Ham  (smoked)            .    .  . 

40.3 

16.1 

38.8 

4.8 

Honey  

18.2 

0.4 

81.2 

0.2 

Mutton  (forequarter)  .  .  . 
Oatmeal 

52-9 
7.3 

J5-3 
16.1 

30.9 
7.2 

67.5 

0.9 

I.O 

Peanuts    

9.2 

25.8 

38.6 

24.4 

2.O 

Potatoes 

78.3 

2.2 

O.I 

18.4 

I.O 

Rice                 

12-3 

8.0 

0.3 

79.0 

0.4 

S  traw  berries 

OO.A 

I.O 

0.6 

7.4 

0.6 

Sugar  (gran.)  

IOO.O 

Tomatoes  

94.3 

0.9 

0.4 

3-9 

°-5 

Walnuts  (English)    

2-5 

16.6 

63-4 

16.1 

1.4 

*  A  fuller  table  can  be  found  in  Bulletin  28  (Revised)  or  Farmers'  Bulletin  142,  United 
States  Department  of  Agriculture,  Office  of  Experiment  Stations. 

Digestive  changes  are  partly  mechanical  and  partly 
chemical.  The  chemical  changes  that  occur  during  diges- 
tion are  due  mainly  to  enzymes,  which  transform  carbo- 
hydrates, fats,  and  proteins  into  substances  which  are  more 
readily  soluble,  diffusible,  and  better  fitted  for  absorption. 
For  example,  pytalin  of  the  saliva  changes  starch  into 
maltose;  other  enzymes  in  the  pancreatic  and  intestinal 


FOOD  AND  NUTRITION  217 

juices  also  act  upon  carbohydrates  and  change  them 
chiefly  into  dextrose.  (See  Part  II,  Exp.  124.)  The 
bulk  of  complex  carbohydrate  material  after  its  change 
into  maltose,  dextrose,  and  similar  sugars  is  absorbed 
by  the  blood.  The  blood  always  contains  a  small  and 
nearly  constant  per  cent  of  dextrose,  but  the  excess  is 
stored  in  the  liver  in  the  form  of  glycogen  (227)  which 
is  given  up  again  gradually  to  the  blood  as  dextrose. 
During  the  circulation  of  the  blood,  dextrose  and  oxygen 
disappear  in  the  muscles  and  eventually  become  carbon 
dioxide  and  water;  much  heat  is  liberated  during  this 
chemical  transformation.  Some  fat  undergoes  complex 
changes,  some  is  transformed  into  glycerin  and  acids 
related  to  palmitic  acid,  partly  in  the  stomach  and 
largely  in  the  intestines,  some  is  stored  in  the  tissues, 
and  some  undoubtedly  undergoes  changes  like  the  dex- 
trose in  the  muscles.  Probably  some  fat  is  converted 
into  carbohydrate.  The  products  of  the  transforma- 
tion of  protein  are  absorbed  by  the  blood  and  thus 
become  available  as  material  for  replacing  worn-out 
tissue  and  contributing  new  tissue  for  the  growth  of  our 
bodies.  Protein  products  are  essentially  body  builders, 
though  some  protein  doubtless  contributes  to  the  forma- 
tion of  carbohydrate  and  fat,  and,  if  these  are  not  supplied 
by  food,  even  performs  their .  functions.  Some  protein 
serves  as  fuel,  but  this  function  belongs  more  especially  to 
carbohydrate  and  fat.  The  end  products  of  the  metabol- 
ism of  protein  are  compounds  of  nitrogen,  chiefly  urea 
(CO(NH2)2)  which  is  excreted  in  solution  from  the  body. 

Water  and  mineral  matter  are  not  foods  in  a  narrow  sense,  because 
they  do  not  build  tissue  or  furnish  energy.  Nevertheless  both  are 
indispensable  for  life  processes.  About  70  per  cent  of  the  weight  of 
the  body  is  water.  It  occurs  in  all  the  tissues  and  keeps  them  soft 


218  CHEMISTRY 

and  pliable.  Water  also  serves  as  a  solvent  and  transporter,  carrying 
juices  and  digested  food  to  all  parts  of  the  body  and  finally  leaving 
the  body  loaded  with  waste  matter.  Some  waste  matter  is  eliminated 
from  the  body  through  the  skin  in  the  sweat.  Mineral  matter  makes 
up  about  4  per  cent  of  the  weight  of  the  body  and  consists  of  com- 
pounds of  the  following  elements  (in  order  of  abundance) :  Calcium, 
phosphorus,  potassium,  sulphur,  sodium,  chlorine,  magnesium,  iron 
(and  minute  quantities  of  iodine,  fluorine,  and  silicon)  (8).  These 
elements  supply  the  material  for  the  rigid  parts  of  the  body.  Thus, 
bones  and  teeth  contain  upwards  of  69  per  cent  of  mineral  matter, 
largely  calcium  phosphate.  They  are  also  essential  constituents  of 
complex  compounds  in  the  tissues  of  important  organs,  e.g.  muscles, 
brain,  and  nerves.  Moreover,  they  furnish  acids,  bases,  salts,  and 
organo-metallic  compounds,  which  give  many  fluids  and  juices  of  the 
body  characteristic  properties.  Illustrations  of  this  function  are  the 
hydrochloric  acid  of  the  gastric  juice  and  the  hematin  of  the  blood 
corpuscles.  Besides  calcium  and  phosphorus  compounds,  sodium 
chloride  is  an  essential  inorganic  compound;  it  is  found  in  many 
parts  of  the  body  and  is  a  necessary  ingredient  of  blood  and  lymph. 
Phosphorus  compounds  are  distributed  in  small  quantities  through- 
out the  body  and  are  as  essential  to  every  living  cell  as  the  protein 
itself,  while  calcium  compounds  are  necessary  for  the  coagulation  of 
the  blood  and  certain  movements  of  the  muscles.  As  a  final  result 
of  metabolism,  inorganic  compounds  are  eliminated  from  the  body 
chiefly  as  chlorides,  sulphates,  and  phosphates  of  sodium,  potassium, 
calcium,  and  magnesium. 

260.  Food  as  a  Source  of  Energy. — The  complex 
mechanical  and  chemical  processes  that  occur  during 
nutrition  not  only  provide  material  for  repairing  and 
building  the  body  but  they  also  yield  energy.  That  is, 
food  is  fuel  in  the  sense  that  it  yields  (i)  heat  energy  to 
keep  the  body  warm  and  maintain  the  temperature  best 
suited  for  the  vital  processes  and  (2)  muscular  and  ner- 
vous energy  to  enable  the  body  to  do  its  work,  e.g.  move 
and  give  motion  to  things  near  it.  Heat  also  results 
from  the  muscular  work  of  the  body,  and  probably  the 


FOOD  AND  NUTRITION  219 

heat  from  both  sources  —  chemical  changes  and  muscular 
work  —  is  sufficient  to  maintain  the  body  temperature. 
In  many  respects  the  human  body  resembles  a  steam 
engine  (18).  Both  get  heat  and  power  from  fuel  —  in 
the  former  case  food,  in  the  latter  coal,  wood,  or  oil;  and 
both  utilize  oxygen  from  the  air  in  this  process.  Waste 
products  likewise  escape  from  both.  But  the  body 
differs  from  the  steam  engine  in  several  important  ways. 
First,  the  body  is  self-building,  self-repairing,  and  self- 
regulating.  Again,  the  body  uses  as  fuel  part  of  the 
material  that  likewise  serves  for  building  and  repairing. 
Moreover,  the  body  is  more  economical  than  the  steam 
engine  in  its  use  of  fuel.  And  finally  the  body  utilizes 
its  fuel  not  merely  for  the  production  of  heat  and  mechan- 
ical energy,  but  also  for  the  exercise  of  its  nervous  system 
and  its  intellectual  and  spiritual  faculties. 

261.  Fuel  Value  of  Food.  — Carbohydrates,  fats,  and 
proteins  may  all   serve  as  fuel  in  the  body.     We  speak 
of  fuel  as  containing  latent  or  potential  energy,  that  is, 
reserve  energy  which  can  be  liberated  and  used  as  needed. 
When    fuel   is    burned,    this    potential    energy    becomes 
kinetic  energy,  that  is,  active,  usable  energy,  e.g.  in  the 
form  of  heat  energy  or  mechanical  energy.     Now  when 
food  is  transformed  chemically  in  the  body,  the  chemical 
change  sooner  or  later  involves  oxidation  or  some  closely 
related  process,  and  as  a  result  the  potential  energy  in 
the  food  becomes  heat  energy  and  muscular  energy  in 
the  body.     Hence  foods  are  said  to  have  "fuel  value." 
And  just  as  various  kinds  of  coal  differ  in  the  amount  of 
heat  liberated  per  ton,  so  various  foods  differ  in  their 
value  as  fuel  for  the  body. 

262.  Determination   of    the    Fuel   Value    of   Food.  - 
The  heat  value  of  food  is  found  by  burning  a  weighed 


220  CHEMISTRY 

quantity  of  the  food  in  a  bomb  calorimeter  and  measuring 
the  amount  of  heat  liberated.  This  amount  of  heat  is 
sometimes  called  the  heat  of  combustion  of  the  substance. 
The  unit  commonly  used  in  measuring  heat  of  combustion 
is  the  large  calorie  (or  Calorie),  that  is,  the  amount  of 
heat  that  would  raise  the  temperature  of  i  kilogram  of 
water  i°  C.  (Compare  164,  174.)  It  is  convenient  to 
think  of  the  large  calorie  as  also  the  amount  of  heat 
energy  equivalent  to  the  mechanical  energy  that  would 
lift  i  ton  1.54  feet  high.  The  bomb  calorimeter  consists 
essentially  of  a  heavy  steel  vessel,  called  a  bomb,  im- 
mersed in  water  in  an  outer  vessel.  A  weighed  amount 
of  the  food  (or  other  substance)  is  first  put  in  the  bomb 
and  a  weighed  amount  of  water  in  the  outer  vessel,  the 
bomb  is  tightly  closed,  oxygen  is  forced  in,  and  the  bomb 
is  then  immersed  in  the  water;  as  soon  as  the  temperature 
of  the  water  is  constant  (or  its  variations  in  temperature 
are  known),  the  substance  inside  the  bomb  is  ignited  by 
electricity,  and  the  rise  in  temperature  of  the  water,  as 
the  substance  burns,  is  accurately  noted. 

Food  as  eaten  is  not  all  digestible  nor  is  the  nutritive 
portion  completely  oxidized  in  the  body,  so  the  heat 
values  of  uneaten  food  obtained  by  the  bomb  calorimeter 
are  not  an  accurate  measure  of  energy  produced  by  the 
bodily  transformation  of  food.  Such  transformations 
are  studied  by  a  respiration  calorimeter.  This  is  a  metal- 
walled  chamber  in  which  a  man  can  live  several  days. 
Experiments  with  different  forms  of  the  respiration  calo- 
rimeter show  that  (i)  the  actual  material  that  is  oxidized 
in  the  body  yields  the  same  amount  of  energy  as  if  it 
were  burned  in  the  bomb  calorimeter,  (2)  when  a  man 
does  no  external  muscular  work,  all  energy  leaves  the 
body  as  heat,  but  when  he  does  muscular  work,  e.g.  lifts 


FOOD  AND  NUTRITION 


221 


weights  or  drives  a  (stationary)  bicycle,  part  of  the 
energy  appears  in  this  external  work  and  the  rest  is  given 
off  as  heat,  and  finally  (3)  the  energy  given  off  as  heat 
when  the  man  rests  or  as  combined  heat  and  mechanical 
energy  when  he  works  equals  the  potential  energy  of  the 
food  material  consumed  in  the  body.  The  third  result 
is  instructive,  for  it  appears  that  the  body  conforms  to 
the  law  of  the  conservation  of  energy,  viz.  energy  can 
be  transformed  without  loss  but  it  cannot  be  destroyed 
or  created.  Fuel  values  calculated  by  elaborate  experi- 
ments with  both  the  bomb  and  respiration  calorimeters 
are  called  physiological  fuel  values.  They  are  the  fuel 
values  of  the  part  of  the  food  that  is  actually  trans- 
formed into  energy  in  the  body,  —  the  real  value  of 
the  food  digested  and  oxidized.  The  physiological  fuel 
values  are  the  ones  usually  meant  when  the  term  fuel 
value  is  used.  It  is  customary  to  state  fuel  value  in 
Calories  per  pound,  i.e.  the  number  of  Calories  furnished 
by  one  pound  of  food.  The  fuel  values  of  the  foods 
whose  composition  is  given  in  258  are  as  follows :  - 

TABLE  or  FUEL  VALUE  OF  FOODS  (CALORIES  PER  POUND) 


Apples 

2QO 

Corn 

4.70 

Peanuts 

2  c6o 

Bananas   .  .  . 

.  .  .  .460 

Eggs  .  .  . 

720 

Potatoes 

381? 

Beans    

.  .  l6o=; 

Grapes  . 

.   4^0 

Rice 

1630 

Beefsteak 

1  1  3O 

Ham 

104.0 

Strawberries 

1  80 

Celery 

8c 

Honey 

i  ^20 

Sugar 

1860 

Cheese    

.    .    .  IQSO 

Mutton 

•  •  I'JO'? 

Tomatoes 

IOC 

Codfish  .... 

.     Tt2< 

Oatmeal 

1860 

Walnuts 

328^ 

263.    Extension   and  Summary  of  Fuel  Value  of  Food. 

-  An  ordinary  mixed  diet  contains  protein,  carbohydrate, 

and  fat.     The  average  proportions  utilized  by  the  body 

are  approximately  protein  92  per  cent,  carbohydrate  97, 


222  CHEMISTRY 

and  fat  95.  This  means,  for  example,  if  53.1  per  cent  of 
a  loaf  of  bread  is  carbohydrate,  the  body  will  appropriate 
as  carbohydrate  approximately  52  per  cent  of  the  weight 
of  the  bread.  It  should  also  be  recalled  that  the  fuel 
value  is  not  the  heat  of  combustion  of  the  uneaten  food, 
but  of  the  food  actually  oxidized  in  the  body.  Inasmuch 
as  protein  is  primarily  a  body  builder  and  carbohydrates 
and  fats  are  heat  producers,  it  is  customary  to  measure 
the  nutritive  value  of  food  by  a  nutritive  ratio.  The 
nutritive  ratio  is  the  ratio  of  digestible  protein  to  diges- 
tible carbohydrates  and  fats  in  a  single  food  or  mixture. 
Such  ratios  indicate  relative  richness  in  nitrogenous  or 
body  building  constituents.  Many  things  about  food  and 
diet  are  summed  up  (as  averages)  in  the  Table  of  Nutri- 
ents, etc.,  on  the  opposite  page. 

264.  Diet  and  Nutritive  Value  of  Food.  —  The  diet 
that  will  build  and  sustain  the  body  of  a  normal  person 
must  contain  the  essential  food  constituents  in  the  correct 
proportions .  Moreover,  the  different  kinds  of  food  to  be 
of  greatest  value  in  a  diet  should  be  properly  prepared 
and  compatibly  mixed.  The  kind  and  quantity  of  the 
foods  that  best  meet  physiological  requirements  vary 
greatly  with  the  age,  activity,  and  environment  of  persons. 
This  subject  has  been  carefully  studied  and  as  a  result 
standard  dietaries  have  been  established.  Formerly, 
standards  were  given  in  terms  of  the  total  nutrients  in 
the  food  eaten.  For  example,  the  proper  amount  of  food 
constituents  needed  per  day  by  a  man  at  moderate  mus- 
cular work  was  given  by  different  authorities  as  protein 
about  120  gm.,  fat  50-65  gm.,  and  carbohydrate  400-531 
gm.  Later  dietaries  are  stated  in  terms  of  protein  and 
energy.  The  table  of  Results  of  Dietary  Studies,  etc.,  on 
the  opposite  page  is  instructive. 


FOOD  AND  NUTRITION 


223 


TABLE  OF  NUTRIENTS,  FUEL  VALUE,  AND  NUTRITIVE  RATIO  OF  COMMON 

FOODS 


$! 

Digestible  Nutrients 

.1.8 

Fuel 

QJ 

Food 

Refuse 

Water 

1| 

Ac.U 

value 

S.8 

"rt  <y 

& 

Pro- 
tein 

Fat 

Carbo- 
hydrate 

Ash 
(rain, 
mat.) 

per 
pound 

I1 

Apples   

25.0 

63.3 

1.2 

•3 

•3 

9-7 

.2 

190 

34.7 

Bananas  

35-o 

48.9 

1.6 

•7 

•4 

12.9 

•5 

260 

19.7 

Beans  (dried)    .  . 

— 

12.6 

7-9 

17-5 

1.6 

57-8 

2.6 

1520 

3-5 

Beef  (loin)  

13-3 

52.5 

1.6 

15-6 

16.6 

— 

•  7 

1025 

2.4 

Bread  (white)  .  . 

— 

35-3 

2.9 

7.8 

1.2 

52.0 

.8 

I2OO 

7.0 

Butter    

— 

II.O 

4.9 

I.O 

80.8 

— 

2.3 

3410 

— 

Cod  (dressed)  .  .  . 

29.9 

58.5 

•5 

10.8 

.2 

— 

.6 

2  2O 

.1 

Eggs  (raw)    

II.  2 

65.5 

i.i 

12.7 

8.8 

— 

•7 

635 

i-7 

Milk              



87.0 

.5 

3.2 

3-8 

5.0 

.5 

310 

4-3 

Mutton  (loin)  .  . 

16.0 

42.0 

2.O 

13-1 

26.9 

•5 

I4IS 

4.6 

Oatmeal  

— 

7-8 

5-i 

14.2 

6.6 

64.9 

1.4 

I800 

5-6 

Potatoes 

2O.O 

62.6 

1.2 

1.5 

.1 

14.0 

.6 

295 

9-5 

Rice   

12.3 

2.9 

6.8 

•3 

77-4 

•3 

1625 

n-5 

Sugar  (gran.)    .  . 



— 



— 

— 

IOO.O 

— 

1860 

— 

Tomatoes    

— 

94-3 

•5 

-7 

A 

3-7 

•4 

95 

6.6 

RESULTS  OF  DIETARY  STUDIES  IN  THE  UNITED  STATES 


Persons  and  Occupation 

Total  Protein 
Eaten  (Gm.) 

Energy  of 
Total   Diet 
(Calories) 

Digested 
Protein 
(Grams) 

Energy 
Utilized 
(Calories) 

Men  at  hard  muscular  work  (Arti- 

sans laborers,  etc.)  

177 

6485 

162 

6000 

Athletes                                

198 

4980 

182 

4"?IO 

Men  at  moderate  muscular  work 

(Farmers,  mechanics,  etc.)  .... 

IOO 

3685 

92 

3425 

Men  not  at  muscular  occupations 

(Business  men,  students,  etc.)  .  . 

106 

3560 

98 

3285 

Men  at  no  muscular  work  (Inmates 

of  institutions)      

86 

2820 

80 

2600 

Very  poor  working  people             .  . 

69 

2275 

64 

2IOO 

224 


CHEMISTRY 


The  nutritive  value  of  food  is  usually  improved  by 
cooking.  If  properly  cooked,  food  is  not  modified  so 
that  the  different  nutrients  fail  ultimately  to  perform 
their  functions  in  building  and  maintaining  the  body. 

If  we  tabulate  different  kinds  of  food  on  the  basis  of 
the  energy  furnished  (by  the  edible  portion)  for  a  given 
sum  of  money,  it  is  seen  that  foods  which  cost  much 
alike  vary  widely  in  their  nutritive  value.  The  accom- 
panying table  shows  the  number  of  Calories  yielded  by 
the  edible  portion  of  various  foods  purchasable  for  one 
dollar  in  each  case:  — 


COMPARATIVE  COST  OF  NUTRIENTS  IN  FOODS 


Food 

Calories 

for  $i 

Food 

Calories 
for$i 

Food 

Calories 
for  $i 

Rolled  oats 
(bulk)    

37,000 

Milk  .      ... 

8,8^5 

Tomatoes 
(canned)    ... 

2,62? 

Sugar 

•21  QOO 

Ham  (smoked) 

7  600 

Halibut 

2,611 

Bread  (white)  . 
Butter 

22,790 

II  26^ 

Beef  (rib)    .... 
Annies 

5,550 

C     CQO 

Oranges  
Oysters    

2,041 

I,C22 

Mutton  (chops) 
Cheese  (cream) 

11,250 

10,833 

Eggs  
Corn  (canned) 

3,735 
3,640 

Strawberries  .  . 
Lobster  

I,4OO 
40O 

EXERCISES 

1.  What  are  carbohydrates?     Why  so  called?     Is  the  term  carbohydrate 
accurate?     Why?     Name  several  carbohydrates. 

2.  Discuss  the  distribution  of  sucrose.     State  its  properties. 

3.  Compare  the  properties  of  dextrose  and  levulose.     How  are  these 
sugars  related  to  sucrose?     Discuss  glucose.     What  is  (a)  grape  sugar  and 
(6)  fruit  sugar? 

4.  Why  are  dextrose  and    levulose  so  named?     What  is  a  test    for 
dextrose  and  similar  sugars? 

6.   Define   and   illustrate  by  means  of  sugars   (a)  hydrolysis  and  (b) 
fermentation. 

6.   Discuss  enzymes. 


FOOD  AND  NUTRITION  225 

7.  How  would  you  show  that  a  leaf  contains  starch? 

8.  Review  relation  of  oxygen  and  carbon  to  plants  and  animals. 

9.  Essay  topics:    (a)  Dextrin,     (b)  Cellulose  —  its   properties   and  de- 
rivatives,     (c)  The  chemistry  of  bread-making,     (d)  The  manufacture  of 
paper,     (e)  Flour.     (/)  Alcohol,     (g)  Glycogen. 

10.  Discuss  the  manufacture  of  ethyl  alcohol. 

11.  What  is  vinegar?     How  is  it  made? 

12.  What  is  cream  of  tartar?     What  is  its  function  in  baking  powder? 

13.  State  the  general  relations  of  fats  to  glycerin.     Name  several  fats. 
What  is  (a)  lard,  (b)  olive  oil,  (c}  butter,  (d)  oleomargarine? 

14.  Discuss  soap  fully. 

16.   Apply  Exercise  14  to  glycerin. 

16.  What  is  meant  by  the  term  protein?    State  some  general  properties 
of  proteins.     Name  the  groups  of  proteins  and  illustrate  each  by  a  familiar 
substance. 

17.  Review  respiration.     What  is  (a)  hemaglobin  and  (b)  hematin? 

18.  Essay  topics:   (a)  Nitrogen  and  life,     (b)  The  blood. 

19.  Define  the  terms  food,  nutrition,  and  nutrients. 

20.  Name  the  three  groups  of  nutrients.     Give  examples  of  each. 

21.  Give  the  composition  of  five  foods. 

22.  Learn  the  composition  and  fuel  value  of  bread,  butter,  and  milk. 

23.  Select  from  the   table  in   258  foods  rich  in  (a)  protein,  (b)  carbo- 
hydrate, and  (c)  fat. 

24.  As  in  Exercise  23,  three  foods  which  make  a  well  balanced  combina- 
tion of  nutrients. 

25.  Essay  topics:  (a)  Digestion,    (b)  Water  and  nutrition,    (c)  Enzymes 
and  digestion,     (d)  Mineral  matter  in  the  body. 

26.  Discuss  the  topics:  (a)  Food  is  a  source  of  energy,    (b)  The  body 
resembles  a  steam  engine. 

27.  State  the  fuel  value  of  five  kinds  of  food. 

28.  Assuming    that   one  square  inch  represents   1000    Calories,  draw 
diagrams  of  the  fuel  values  of  ten  foods  from  the  table  in  262. 

29.  Define  the  term  nutritive  ratio.     Give  several  nutritive  ratios.     What 
is  the  nutritive  ratio  of  bread  and  of  milk? 

PROBLEMS 

1.  How  much  sodium  hydroxide  (95  per  cent  pure)  is  needed  to  make 
a  metric  ton  of  soap  (sodium  palmitate)? 

2.  Suppose  35  gm.  of  absolute  ethyl  alcohol  burn  to  water  and  carbon 
dioxide.     What  volume  of  oxygen  at  20°  C.  and  765  mm.  is  needed? 

3.  If  75  gm.  of  dextrose  were  transformed  into  carbon  and  water,  how 
much  of  each  would  be  produced? 


CHAPTER  XVIII 

SULPHUR  — SULPHIDES  — SULPHUR  OXIDES,  ACIDS, 
AND   SALTS 

265.  Occurrence.  —  Large  quantities  of  free   sulphur 
are  frequently  found  in  regions  where   there  are  active 
or  extinct  volcanoes,  as  in  Japan,  Spain,  and  Mexico. 
There  are  other  deposits,  as  in  Sicily,  which  were  doubt- 
less formed  by  the  action  of  animal  or  vegetable  matter 
on    sulphur    compounds,    especially    calcium    sulphate. 

Combined  sulphur  is  found  as  sulphides  and  sulphates. 
Some  important  metallic  sulphides  are  lead  sulphide 
(galena,  PbS),  zinc  sulphide  (sphalerite  or  zinc  blende, 
ZnS),  mercuric  sulphide  (cinnabar,  HgS),  and  copper 
sulphide  (chalcocite,  Cu2S,  and  chalcopyrite,  CuFeS2). 
Iron  sulphide  (iron  pyrites,  FeS2)  is  an  abundant  and 
important  compound.  The  most  abundant  sulphates  are 
calcium  sulphate  (gypsum,  CaSO^H^O)  and  barium 
sulphate  (BaS04).  Volcanic  gases  often  contain  sulphur 
dioxide  (S02),  and  hydrogen  sulphide  (H2S)  is  found  in 
the  water  of  sulphur  springs.  Sulphur  is  a  constituent 
of  certain  organic  compounds  present  in  onions,  horse- 
radish, mustard,  and  eggs. 

266.  Source.  —  Sicily,   until  the  last  few  years,   fur- 
nished most  of  the  sulphur.     Some  is  also  provided  by 
Japan,  Spain,  Iceland,  and  Mexico.     Enough  sulphur  for 
domestic  and  some  foreign  uses  is  now  readily  obtained 
from  the  vast  deposits  in  Louisiana. 


SULPHUR 


227 


267.  Extraction  and  Purification.  —  Sicilian  sulphur  contains  a 
large  per  cent  of   rock  and   earth.      It  is  extracted  by  igniting  a 
pile  of  the  ore.      The  heat  produced  by  the  combustion  of   some 
of  the  sulphur  melts  the  rest,  which  drains  away  from  the  impuri- 
ties.     The  crude  sulphur 

thus  obtained  is  purified 
in  an  apparatus  like  that 
shown  in  Fig.  66.  The 
crude  sulphur  is  melted 
in  the  reservoir  B  and 
distilled  in  the  iron  vessel 
A;  the  vapors  pass  into 
the  large  brick  chamber 
and  condense  upon  the 
cold  walls  as  a  fine  pow- 
der, called  flowers  of  sul- 
phur. As  the  operation  jig.  66.  —  Apparatus  for  Purifying  Sulphur, 
continues  the  walls  be- 
come hot,  and  the  sulphur  collects  on  the  floor  as  a  liquid;  this  is 
drawn  off  through  C  into  cylindrical  molds  and  allowed  to  cool. 
These  sticks  are  called  roll  sulphur  or  brimstone. 

In  Louisiana  the  sulphur  is  obtained  in  a  pure  state  by  forcing 
very  hot  water  through  pipes  down  upon  the  deep  beds;  the  melted 
sulphur  rises  part  way  in  another  pipe,  and  is  pumped  to  the  sur- 
face by  compressed  air.  The  sulphur  wells,  as  they  are  called, 
are  very  powerful,  a  single  well  often  pumping  500  tons  daily. 
The  liquid  sulphur  flows  into  large  wooden  bins  where  it  solidifies. 
Since  Louisiana  sulphur  contains  over  99  per  cent  of  the  element, 
no  purification  is  necessary. 

268.  Properties.  —  Sulphur  is  a  yellow,  brittle  solid 
which  sometimes  has  a  faint  odor.     It  is  insoluble  in 
water,  but  most  varieties  dissolve  in  carbon  disulphide. 
Sulphur  does  not  conduct  electricity;  nor  does  it  conduct 
heat  well,  even  the  warmth  of   the   hand   causing  it   to 
crackle  and  break.     The  specific  gravity  of  the  solid  is 
about  2.     At  the  lowest  temperature  at  which  sulphur 
can  be  vaporized,  a  molecule  contains  eight  atoms  (S8), 


228  CHEMISTRY 

while  at  800°  C.  and  higher  it  contains  two  atoms  (S2). 
When  heated,  sulphur  melts  to  a  thin,  amber-colored 
liquid.  As  the  temperature  is  raised,  the  liquid  darkens 
and  thickens,  until  at  about  230°  C.  it  is  black  and  too 
thick  to  be  poured  from  the  vessel.  Heated  still  higher, 
the  color  remains  black  but  the  mass  becomes  thin,  and 
finally  at  about  448°  C.  the  liquid  boils  and  turns  into 
a  yellowish  vapor. 

Sulphur  combines  readily  and  directly  with  many  ele- 
ments, thereby  forming  sulphides.  Thus,  it  ignites  readily 
in  the  air  and  burns  with  a  pale  blue  flame,  forming  sul- 
phur dioxide  gas  (SC^) ;  if  burned  in  oxygen,  a  little 
sulphur  trioxide  (SO3)  is  also  formed.  With  metals  the 
reaction  is  often  accompanied  with  much  heat  and  light. 
(See  Part  II,  Exps.  130,  133.) 

269.  Different  Modifications  of  Sulphur.  —  Sulphur 
exists  in  several  different  modifications,  which  fall  into 
two  classes  —  crystallized  and  amorphous. 

Crystallized  sulphur  includes  the  varieties  that  produce 
crystals  belonging  to  the  ortho- 
rhombic  and  monoclinic  systems. 
Orthorhombic  crystals  of  sulphur 
are  deposited  by  a  slowly  evapora- 
ting solution  of  carbon  disulphide 
(Fig.  67).  Crystallized  native  sul- 

Fig.  67. — •  Orthorhombic          ,         .         ..11         i  •        TV/T  v    • 

Crystals  of  Sulphur.         Phur  1S  Orthorhombic.    Monoclinic 
crystals  are  deposited  from  slowly 

cooling  molten  sulphur.  By  melting  sulphur  in  a  crucible 
and  pouring  off  the  excess  of  liquid  as  soon  as  crys- 
tals shoot  out  from  the  walls  near  the  surface,  the 
interior  of  the  crucible  is  found  to  be  lined  with  long, 
dark  yellow,  shining  needles  (Fig.  68).  They  are  mono- 
clinic  crystals  of  sulphur. 


SULPHIDES 


229 


Fig.  68. — Section  of  a  Cru- 
cible Showing  Monoclinic 
Crystals  of  Sulphur. 


Amorphous    sulphur    is    formed    by    pouring     boiling 
sulphur  into  cold  water.     It  is  a  tough,  rubberlike,  am- 
ber-colored solid,  which  is  called  plastic  sulphur.     It  is 
insoluble  in  carbon  disulphide.    It 
soon  becomes  hard,  brittle,  and  yel- 
low, and   after  a  time  changes  in 
part  into  orthorhombic  crystals.    If 
these  are  dissolved  out  by  carbon 
disulphide,     a    permanent    amor- 
phous form  remains.    (See  Part  II, 
Exps.  131,  132.) 

270.  Uses.  —  Large    quantities 
of  sulphur  are  used  in  making  sul- 
phuric acid  and  other  sulphur  com- 
pounds.    Considerable  is  also  used 
in  the  manufacture  of  gunpowder, 

fireworks,  matches,  hard  and  soft  rubber,  and  mixtures 
for  killing  insect  pests.  These  mixtures,  e.g.  lime-sul- 
phur spray,  contain  unstable  sulphur  compounds,  which 
readily  decompose  and  liberate  sulphur  upon  the  injuri- 
ous insect.  Sulphur  itself  is  also  used  as  an  insecticide, 
especially  for  killing  Phylloxera  —  an  insect  which  de- 
stroys grapevines. 

Sulphides 

271.  Hydrogen  Sulphide,  H2S,  is  a  gaseous  compound 
of  sulphur  and  hydrogen.     It  is  sometimes  called  sul- 
phuretted hydrogen.     It  occurs  in  the  waters  of  sulphur 
springs.     It  is  found  in  the  air,   especially  near  sewers 
and  cesspools,   since  it  is  one  product  of  the   decay  of 
animal  substances,  many  of  which  contain  sulphur.     For 
example,  the  albumin  in  the  white  part  of  eggs  is  rich 
in  sulphur,  and  when  eggs  decay,  hydrogen  sulphide  is 
readily  formed.     (See  Part  II,  Exp.  103C.) 


230  CHEMISTRY 

272.  Preparation  of  Hydrogen  Sulphide.  --  The  gas  is 
prepared  in  the  laboratory  by  the  interaction  of  dilute 
acids  and  metallic  sulphides,  usually  by  hydrochloric  acid 
and  ferrous  sulphide.     The  equation  is:  - 

FeS     +       2HC1      =     H2S     +     FeCl2 

Iron  Hydrochloric         Hydrogen  Iron 

Sulphide  Acid  Sulphide  Chloride 

273.  Properties   of   Hydrogen    Sulphide.  —  Hydrogen 
sulphide  gas  is  colorless  and  has  the  odor  of  rotten  eggs. 
It  is  poisonous.     A  little,  even  if  diluted  with  air,  often 
produces  headache  and  nausea,  and  a  large  quantity  of 
the  gas  may  prove  fatal.     One  volume  of  water  dissolves 
about  three  volumes  of  hydrogen  sulphide  gas  at  ordinary 
temperatures.      The   solution   is   often    called   hydrogen 
sulphide  water,  and  can  be  used  instead  of  the  gas  in 
many  chemical  experiments;   it  has  a  weak  acid  reaction, 
and  hence  hydrogen  sulphide  is  sometimes  called  hydro- 
sulphuric  acid.    In  terms  of  the  ionic  theory  the  solution 
contains  relatively  few  hydrogen  ions.       (See   Part   II, 
Exp.  138.) 

Hydrogen  sulphide  gas  is  inflammable  and  burns  with 
a    bluish    flame,    forming    sulphur    dioxide    and    water, 

2H2S    +    302    =    2S02    +    2H20 

Hydrogen        Oxygen        Sulphur  Water 

Sulphide  Dioxide 

If  the  supply  of  air  is  insufficient,  combustion  is  incom- 
plete, sulphur  and  water  being  formed,  thus:  - 

2H2S     +     02     =     28     +     2H20 

Hydrogen        Oxygen        Sulphur  Water 

Sulphide 

Hydrogen  sulphide  reduces  nitric  acid  and  sulphuric  acid, 
the  equation  for  the  latter  reaction  being:  - 


SULPHIDES  231 

H2S      +  H2S04    =    S02    +      S       +  H20 

Hydrogen        Sulphuric        Sulphur         Sulphur        Water 
Sulphide  Acid  Dioxide 

274.  Sulphides  may  be  regarded  as  salts  of  the  weak 
acid,  hydrogen  sulphide,  though  they  are  not  always 
prepared  directly  from  hydrogen  sulphide.  They  can  be 
produced  by  the  direct  union  of  sulphur  and  metals,  as 
in  the  case  of  iron  and  copper  sulphides  previously  men- 
tioned, or  by  exposing  metals  to  the  moist  gas.  Some 
are  prepared  by  heating,  or  merely  mixing,  an  hydroxide 
with  an  excess  of  sulphur;  to  this  class  belong  the  com- 
plex, unstable  compounds  like  ammonium  polysulphide  and 
calcium  persulphide.  A  more  common  way  is  to  pre- 
cipitate them  by  passing  the  gas  into  solutions  of  metallic 
compounds,  or  by  adding  hydrogen  sulphide  water.  Cop- 
per, tin,  lead,  and  silver  react  readily  with  hydrogen 
sulphide  and  are  rapidly  tarnished  by  exposure  to  the 
gas.  Silverware,  on  this  account,  turns  brown  or  black, 
especially  in  houses  heated  by  coal  and  lighted  by  coal 
gas.  The  brown  silver  sulphide  also  coats  silver  spoons 
which  are  put  into  mustard,  eggs,  and  some  vegetables, 
such  as  cauliflower.  Lead  compounds  are  blackened  by 
this  gas,  owing  to  the  formation  of  lead  sulphide,  thus :  - 

PbO     +  H2S     =  PbS     +     H2O 

Lead  Oxide        Hydrogen  Sulphide        Lead  Sulphide        Water 

For  this  reason  houses  painted  with  " white  lead"  paint 
often  become  dark,  and,  similarly,  oil  paintings  are  dis- 
colored. The  blackening  of  paper  moistened  with  a  solu- 
tion of  lead  nitrate  or  acetate  is  the  customary  test  for 
hydrogen  sulphide.  (See  Part  II,  Exp.  139.) 

Many  sulphides  are  black,  one  (zinc  sulphide)  is  white,  and 
several  have  a  characteristic  color.  Thus,  arsenic  sulphide  is 
pale  yellow,  cadmium  sulphide  is  golden  yellow,  manganese  sulphide 


232  CHEMISTRY 

is  flesh  colored,  and  antimony  sulphide  is  orange  red.      The  color 
often  affords  a  ready  means  of  detecting  each  sulphide. 

Sulphur  Oxides,  Acids,  and  Salts 

275.  Sulphur   Dioxide,  S02,  is  the  usual  product  of 
the  combustion  of  sulphur  and  sulphur  compounds. 

276.  Formation  and  Preparation  of  Sulphur  Dioxide. 
—  When  sulphur  burns  in  air  (or  oxygen) ,  sulphur  di- 
oxide is  formed,  thus:- 

S     +      O2       =        SO2 

Sulphur        Oxygen        Sulphur  Dioxide 

It  is  also  formed  by  roasting  metallic  sulphides,  especially 
iron  disulphide  (iron  pyrites,  FeS2)  in  air,  thus:  - 

4FeS2     +      nOz  8SO2     +        2Fe203 

Iron  Disulphide         Oxygen         Sulphur  Dioxide         Iron  Oxide 

Both  of  these  reactions  are  utilized  on  a  large  scale  in  the 
manufacture  of  sulphuric  acid  (284). 

277.  Preparation  of  Sulphur  Dioxide.  --  Two  methods 
of  preparation  are  used  in  the  laboratory,     (i)  If  copper 
and  concentrated  sulphuric  acid  are  heated,  a  series  of 
complex  changes  results  finally  in  the  evolution  of  sulphur 
dioxide.     The  equation  may  be  written:  - 

Cu     +     2H2SO4     =     SO2     +       CuSO4     +     2H2O 

Copper        Sulphuric  Acid        Sulphur        Copper  Sulphate 

Dioxide 

(2)  Dilute  sulphuric  (or  hydrochloric)  acid  dropped  upon 
a  sulphite  yields  sulphur  dioxide,  thus :  - 

Na2SO3     +     H2SO4     =     SO2     +     Na2SO4     +     H2O 

Sodium  Sulphuric  Sulphur  Sodium 

Sulphite  Acid  Dioxide  Sulphate 

The  sulphite  method  is  safer  and  more  convenient,  espe- 
cially for  liberating  a  steady  current  of  the  gas. 


SULPHUR   OXIDES,  ACIDS,  AND    SALTS       233 

278.  Properties    of    Sulphur    Dioxide.  —  Sulphur    di- 
oxide gas  has  no  color.      Its  odor  is  suffocating,  being 
the  well-known  odor  associated  with  burning  sulphur.     A 
liter  of  sulphur  dioxide   at  o°   C.   and  760  mm.  weighs 
2.927  gm.     The  gas  is  very  soluble  in  water;   at  ordinary 
temperatures  one  volume  of  water  dissolves  about  forty 
volumes  of  gas.     This  solution  contains  sulphurous  acid 
(280).     (See  Part  II,  Exps.  134,  140.) 

Sulphur  dioxide  will  not  burn  in  air,  nor  will  it  support 
ordinary  combustion.  Moist  sulphur  dioxide  bleaches 
organic  coloring  matters.  A  red  or  a  purple  flower  loses 
color  in  it.  Silk,  hair,  straw,  wool,  and  other  delicate 
substances,  which  would  be  injured  by  chlorine,  are 
sometimes  bleached  by  sulphur  dioxide.  Sulphur  dioxide 
and  sulphurous  acid  are  injurious  substances.  In  locali- 
ties where  considerable  sulphur  dioxide  escapes,  e.g.  in 
factory  and  mining  towns,  trees  are  killed  and  stone 
buildings  are  slowly  disintegrated;  sometimes  plants  are 
killed  in  rooms  where  sulphur  dioxide  is  formed  by  burn- 
ing impure  illuminating  gas. 

279.  Uses  of  Sulphur  Dioxide.  —  Large  quantities  of 
sulphur  dioxide  are  used  in  the  manufacture  of  sulphuric 
acid.     The  gas  is  also  used  as  a  disinfectant  and  pre- 
servative, in  paper  making,  in  tanning,  in  refining  sugar, 
and  in  making  acid  sulphites. 

280.  Sulphurous  Acid.  •*-  Sulphurous  acid   (H2S03)    is  prepared 
by  dissolving  sulphur  dioxide  in  water.      Sulphur  dioxide  is,  there- 
fore, sulphurous  anhydride  (187).      The  simplest  equation  express- 
ing  this  fact  is:  - 


Sft          +  ^Q   =        H!S()J 

Sulphur  Dioxide         Water          Sulphurous  Acid 

The  acid  is  unstable  and  decomposes  readily  into  sulphur  dioxide 
and  water;  solutions  of  sulphurous  acid  smell  strongly  of  sulphur 
dioxide.  It  is  also  oxidized  readily;  solutions  of  the  acid,  if  ex- 


234  CHEMISTRY 

posed  to  air  for  a  short  time,  give  a  test  for  sulphuric  acid  which  is 
formed  by  the  combining  of  the  sulphurous  acid  with  oxygen  from 
the  air.  Oxidizing  agents,  such  as  potassium  permanganate,  pro- 
duce this  change  quickly. 

281.  Sulphites.  —  Sulphurous  acid  forms  two  classes  of  salts  — 
the  normal  and  acid  sulphites.      They  yield  sulphur  dioxide  when 
treated  with  acids.      Acid  sodium  sulphite  (HNaSO3),  often  called 
bisulphite  of  soda,  is  the  antichlor  used  to  remove  the  excess  of 
chlorine  from  bleached   cotton  cloth.      It  is  also  used  in  brewing, 
tanning,   and    in    making    starch,    sugar,   and    paper.       Acid    cal- 
cium sulphite   (CaH2(SO3)2)   is  extensively   used  to  dissolve  resin- 
ous and  non-fibrous  matter  in  one  process  of  making  paper  from 
wood  (231). 

282.  Sulphur   Trioxide,  S03,  is  produced   to  a  slight 
extent  when    sulphur   burns   in   air    or   in    oxygen    and 
causes  the  white  fumes  often  seen  during  the  combustion. 
The  combining  of  sulphur  dioxide  and  oxygen   can  be 
hastened  by  passing  a  mixture  of  sulphur  dioxide  and 
oxygen   (or  air)    over  hot  platinum   or  asbestos  coated 
with  platinum.     Other  substances  also  hasten  the  chemi- 
cal change.     The  platinum,  as  far  as  we  know,  does  not 
participate  in  this   reaction.      It    hastens   a  very   slow 
chemical  reaction;  its  function  is  sometimes  said  to  be 
like  that  of  oil  which  assists  the  movement  of  machinery. 
Substances  which  hasten  a  chemical  reaction  but  are  un- 
changed at  the  end  of  the  process  are  called  catalyzers  or 
catalytic  agents.     Such  a  chemical  reaction  is  said  to  be 
due  to  catalysis  or  catalytic  action. 

Sulphur  trioxide  below  15°  C.  is  a  white  crystalline 
solid;  at  15°  C.  it  melts  and  at  46°  C.  it  boils.  Another 
solid  form,  silk-like  in  luster  and  appearance,  is  known. 
When  exposed  to  moist  air,  it  fumes  strongly,  forming 
sulphuric  acid;  and  when  dropped  into  water  it  dissolves 
with  a  hissing  sound  and  evolution  of  heat. 


SULPHUR  OXIDES,  ACIDS,  AND  SALTS        235 

283.  Sulphuric  Acid,  H2SO4,  is  the  best  known  and 
most  largely  used  acid.     Enormous  quantities  are  manu- 
factured, and  the  acid  is  an  indispensable  substance  not 
only  in  many  important  industries  but  also  in  a  vast 
number  of  chemical  processes. 

284.  Manufacture  of  Sulphuric  Acid.  —  Sulphuric  acid 
is  manufactured  by  two  processes,  known  as  the  lead 
chamber  process  and  the  contact  process. 

The  lead  chamber  process  consists  in  introducing  sul- 
phur dioxide,  air,  steam,  and  oxides  of  nitrogen  into  large 
lead  chambers.  Under  suitable  conditions  these  gases 
react  and  produce  sulphuric  acid,  which  collects  on  the 
floors  of  the  lead  chambers. 

The  chemical  changes  involved  in  this  process  of  manu- 
facturing sulphuric  acid  are  complex  and  variable.  The 
main  reactions  may  be  represented  thus:  — 

2S02  +  NO  +  N02  +  02  +  H20  =  2HS04NO 

Nitrosyl-sulphuric 
Acid 

2HS04NO  +  H2O  =  2H2S04  +  NO  +  NO2 

(The  formula  of  nitrosyl-sulphuric  acid  is  sometimes 
written  HSO3NO2,  HOSO2N02,  and  HOSO2NOO.)  An 
examination  of  these  equations  shows  that  the  oxides  of 
nitrogen  play  an  important  part  in  the  manufacture  of 
sulphuric  acid.  The  oxides  of  nitrogen  in  the  second 
equation  are  the  same  in  quantity  as  that  needed  for  the 
first  equation.  Apparently  only  a  small  quantity  of  these 
oxides  would  be  needed  to  form  an  unlimited  quantity 
of  sulphuric  acid.  As  a  matter  of  fact,  some  of  the  oxides 
escape  and  must  be  replaced;  the  loss  is  compensated  by 
putting  nitric  acid  into  the  top  of  the  Glover  tower  or 
injecting  it  into  the  chambers. 


236 


CHEMISTRY 


285.  The  Con- 
struction and  Oper- 
ation of  a  Chamber 
Acid  Plant  can  be 
grasped  by  a  study 
of  Fig.  69.  The 
plant  consists  of 
three  main  parts 
—  (a)  the  furnace 
for  producing  sul- 
phur dioxide,  (b) 
the  lead  chambers 
together  with  the 
Glover  and  Gay- 
Lussac  towers  for 
changing  the  gases 
into  sulphuric  acid, 
and  (c)  the  concen- 
trating apparatus. 
The  manufacturing 
operation  is  some- 
what as  follows : 
(i)  Sulphur  or  iron 
disulphide  (FeS2) 
is  burned  in  a  fur- 
nace constructed 
so  that  enough  air 
passes  over  the 
burning  mass  to 
change  the  sulphur 
into  sulphur  diox- 
ide and  to  furnish 
the  proper  amount 
of  oxygen  for  later 
changes.  In  some 
works  the  furnace 
is  provided  with 


Fig.  69. — Apparatus  for  Chamber  Process. 


SULPHUR  OXIDES,  ACIDS,  AND  SALTS        237 

"  niter  pots "  containing  a  mixture  of  sodium  nitrate  and  sul- 
phuric acid;  the  nitric  acid  vapors  which  are  liberated  are  one 
source  of  the  oxides  of  nitrogen.  (2)  The  mixture  of  sulphur 
dioxide,  oxides  of  nitrogen,  and  air  passes  from  the  furnace  into 
the  bottom  of  the  Glover  tower.  This  is  a  tall  tower  filled  with 
small  acid-resisting  stones  or  pieces  of  special  shaped  earthenware 
over  which  flow  two  streams  of  sulphuric  acid,  one  dilute  and  the 
other  containing  nitrogen  dioxide  (obtained  from  the  Gay-Lussac 
tower  acid).  These  acids  not  only  cool  the  ascending  gases,  but 
the  dilute  acid  is  deprived  of  water  and  the  Gay-Lussac  tower  acid 
of  its  dissolved  nitrogen  dioxide.  Hence,  concentrated  acid  flows 
out  of  the  bottom  of  the  Glover  tower  into  the  "  strong  acid " 
blower  from  which  it  is  forced  to  the  top  of  the  Gay-Lussac  tower 
for  use  as  explained  below;  while  from  the  top  of  the  Glover  tower 
sulphur  dioxide,  nitrogen  dioxide,  and  air  pass  on  into  the  first 
lead  chamber.  Here  steam  is  introduced,  and  often  nitric  acid 
vapor.  The  main  chemical  changes  by  which  the  sulphuric  acid 
is  formed  occur  in  this  and  the  second  chamber.  These  chambers 
are  huge  boxes  often  having  a  total  capacity  of  150,000  cubic  feet; 
the  walls  and  floors  are  of  sheet  lead  supported  on  a  wooden  frame- 
work, lead  being  a  metal  which  is  only  slightly  attacked  by  the 
chamber  acid.  The  nitrogen  (from  the  original  air)  and  the  unused 
gases  pass  on  into  the  bottom  of  the  Gay-Lussac  tower.  This 
tower  is  filled  with  coke  over  which  flows  concentrated  sulphuric 
acid  (obtained  from  the  Glover  tower),  which  absorbs  the  unused 
nitrogen  dioxide.  The  "  nitrose  "  acid  flows  from  the  bottom  of 
the  Gay-Lussac  tower  into  a  blower,  whence  it  is  forced  to  the  top 
of  the  Glover  tower,  where  as  stated  above  the  nitrogen  dioxide  is 
liberated.  At  the  end  of  the  plant  there  is  a  very  tall  chimney, 
which  serves  as  an  exit  for  unused  gases  (such  as  nitrogen)  and 
also  maintains  a  draft  strong  enough  to  draw  the  gases  through 
the  chambers  and  towers.  (3)  The  acid  that  is  produced  in  the 
chambers  and  drawn  off  at  intervals  contains  about  67  per  cent  of 
the  compound  H2SO4.  For  some  uses,  e.  g.  the  manufacture  of 
phosphate  fertilizers,  this  acid  needs  no  further  treatment.  Much 
of  the  acid,  however,  is  concentrated  by  evaporation.  It  is  first 
heated  in  lead-lined  pans  until  the  concentration  is  about  80  per 
cent,  and  finally  in  platinum,  cast  iron,  or  fused  quartz  vessels 
until  the  acid  contains  about  95  per  cent  of  H2SO4. 


238  CHEMISTRY 

286.  Manufacture  of  Sulphuric  Acid  by  the  Contact 
Process. -- This    process    consists    in    bringing    sulphur 
dioxide  and  air,  well  purified  and  heated  to  about  400°  C.,  in 
contact  with  a  catalytic  agent,  which  is  usually  platinum. 
The  sulphur  dioxide  is  thereby  oxidized  to  sulphur  tri- 
oxide,  thus :  - 

2SO2       +      O2  2SO3 

Sulphur  Dioxide        Oxygen        Sulphur  Trioxide 

The  resulting  sulphur  trioxide  is  conducted  into  sulphuric 
acid  containing  a  little  water  with  which  the  sulphur 
trioxide  combines,  thus  producing  very  concentrated 
sulphuric  acid,  which  can  be  used  unchanged  or  diluted 
to  any  strength,  if  desired.  The  equation  for  the  final 
reaction  is:  - 

SO3      +      H2O     =     H2SO4 

Sulphur  Trioxide        Water        Sulphuric  Acid 

287.  The  Construction  and  Operation  of  a  Contact  Acid  Plant  can 
be  best  understood  by  the  sketch  shown  in  Fig.  70.      The  blower  A 
forces  air  into  the  burner  B,  where  the  sulphur  dioxide  is  formed 
by  burning  iron  pyrites  (FeS2)  or  sulphur.      From  the  burner  the 
gases   pass  into    the  dust   chamber  C,  where  they  are  freed  from 
sulphur  dust  and  other  solid  impurities;   this  is  an  important  step, 
for  dust   reduces   the   transforming  power  of   the  catalytic  agent. 
The  gases,   cooled    by   the    pipe    D,   are   further  cleaned    in    the 
scrubbers,  which  contain  coke  wet  with  water  (E)  and   with  sul- 
phuric acid  (F).      The  next   step  is  the  removal   of  arsenic  com- 
pounds in  the  purifier  G.      During  the  combustion  of  iron  pyrites, 
arsenic    compounds    are     liberated;     traces    of     such    compounds 
"  poison "   the  platinum  used  as  a  catalytic  agent   and  stop  the 
formation  of  sulphur  trioxide.      The  purified  gases  (mainly  sulphur 
dioxide)  then  enter  the  mixer  and  heater  H.      Here  a  large  excess 
of   air  is  introduced  from   the   blower  arid   the  whole  mixture  is 
heated  to  about  400°  C.      The  purified  and  heated  mixture  of  sul- 


SULPHUR  OXIDES,  ACIDS,  AND  SALTS        239 


phur  dioxide  and  air  next  passes 
into  the  contact  chamber  I.  Here 
the  gases  come  in  contact  with  the 
catalytic  agent  and  form  sulphur 
trioxide.  The  catalytic  agent,  if 
platinum,  consists  of  asbestos  fibers 
coated  with  a  very  thin  layer  of 
metallic  platinum  and  is  spread  out 
on  plates  or  mixed  with  porous  ma- 
terial in  order  to  provide  a  large 
contact  surface.  The  final  step  is 
the  transformation  of  the  sulphur 
trioxide  into  sulphuric  acid  by  com- 
bining with  water.  Since  all  the 
sulphur  trioxide  formed  does  not 
combine  if  absorbed  in  pure  water, 
the  trioxide  is  passed  into  the  ab- 
sorber J,  which  is  partly  filled  with 
sulphuric  acid  containing  only  2  to 
3  per  cent  of  water.  In  this  liquid, 
combination  takes  place  readily, 
water  being  added  to  maintain 
the  required  concentration  in  the 
absorber. 

288.  Properties  of  Sul- 
phuric Acid.  —  Sulphuric  acid 
is  an  oily  liquid,  colorless 
when  pure,  but  often  brown 
from  the  presence  of  charred 
organic  matter,  such  as  dust 
and  straw.  The  specific 
gravity  of  the  commercial 
acid  is  about  1.83.  It  boils 
at  about  338°  C.,  decompos- 
ing to  some  extent  into  water 
and  dense,  white,  suffocating 


240  CHEMISTRY 

fumes  of  sulphur  trioxide;    decomposition  begins  at  a 
much  lower  temperature  (ioo°  — 150°  C.). 

When  sulphuric  acid  is  mixed  with  water,  much  heat 
is  evolved.  The  acid  should  always  be  poured  into  the 
water  and  the  mixture  should  be  stirred,  otherwise  the 
intense  heat  may  crack  the  vessel  or  spatter  the  hot 
acid.  The  tendency  to  absorb  water  is  shown  in  many 
ways.  The  concentrated  acid  absorbs  moisture  from  the 
air  and  from  gases  passed  through  it.  It  is  often  used  in 
the  laboratory  to  dry  gases.  Organic  substances,  such 
as  wood,  paper,  sugar,  starch,  and  cotton,  are  charred  by 
sulphuric  acid.  Such  compounds  contain  hydrogen  and 
oxygen  in  the  ratio  to  form  water;  these  two  elements 
are  abstracted  and  carbon  alone  remains.  Sulphuric  acid 
also  disintegrates  the  flesh,  often  causing  serious  burns, 
and  if  accidentally  spilled  on  the  hands  or  spattered  on 
the  face  should  be  washed  off  immediately.  The  inter- 
action of  sulphuric  acid  and  metals  varies.  With  many 
metals,  dilute  sulphuric  acid  forms  hydrogen  and  the 
corresponding  metallic  sulphate.  Thus,  hydrogen  is  usu- 
ally prepared  in  the  laboratory  and  often  industrially 
from  zinc  and  sulphuric  acid.  Concentrated  acid,  espe- 
cially if  hot,  converts  most  metals  into  an  oxide,  which  is 
then  transformed  by  the  acid  into  the  corresponding  sul- 
phate. Thus,  the  equations  expressing  the  reactions  with 
copper  are :  — 

Cu    +      H2SO4  CuO      +         SO2          +  H2O 

Copper        Sulphuric  Acid        Copper  Oxide        Sulphur  Dioxide       Water 

CuO       +      H2SO4  CuSO4       +  H2O 

Copper  Oxide        Sulphuric  Acid        Copper  Sulphate        Water 

Iron  is  the  only  common  metal  that  is  not  readily  attacked 
by  the  concentrated  acid,  and  advantage  is  taken  of  this 


SULPHUR  OXIDES,  ACIDS,  AND  SALTS       241 

property  in  transporting  acid  in  bulk  in  iron  tank  cars. 
Sulphuric  acid  unites  with  ammonia  (NH3)  to  form 
ammonium  sulphate  ((NH4)2S04).  By  interaction  with 
salts  the  corresponding  acid  is  readily  liberated  if  the  acid 
boils  at  a  lower  point  than  sulphuric  acid;  this  reaction 
is  widely  used  to  prepare  hydrogen  chloride  from  chlorides 
and  nitric  acid  from  nitrates.  Dilute  solutions  of  sul- 
phuric acid  contain  an  abundance  of  hydrogen  ions  (H+) 
and  sulphate  ions  (SO4  ).  (See  Part  II,  Exp.  135.) 

The  uses  of  sulphuric  acid  are  numerous.  It  is  one 
of  the  most  important  substances.  Directly  or  indirectly 
it  is  used  in  hundreds  of  industries  upon  which  the  com- 
fort, prosperity,  and  progress  of  mankind  depend.  It  is 
used  in  the  manufacture  of  many  acids.  It  is  essential  in 
one  process  for  the  manufacture  of  sodium  carbonate, 
which  has  many  uses.  Enormous  quantities  are  consumed 
in  making  fertilizers,  alum  and  other  sulphates,  nitro- 
glycerin,  glucose,  dyestuffs,  and  in  various  parts  of  such 
fundamental  industries  as  dyeing,  bleaching,  metal  clean- 
ing, refining,  and  metallurgy. 

289.  Sulphates.  —  Sulphuric  acid  is  dibasic  and  forms 
two  classes  of  salts,  —  the  normal  sulphates,  such  as 
sodium  sulphate,  Na2SO4,  and  the  acid  sulphates,  such 
as  acid  sodium  sulphate,  HNaSO4.  Most  sulphates  are 
soluble  in  water,  only  the  sulphates  of  barium,  strontium, 
and  lead  being  insoluble,  while  calcium  sulphate  is  slightly 
soluble.  Important  sulphates  are  calcium  sulphate  (gyp- 
sum, CaS04.2H2O),  barium  sulphate  (barite,  barytes, 
heavy  spar,  BaSO4),  zinc  sulphate  (ZnSO4,  and  white 
vitriol,  ZnSO4.yH2O),  copper  sulphate  (CuSO4,  and  blue 
vitriol  or  blue  stone,  CuS04.5H2O),  iron  sulphate  (FeSO4, 
and  green  vitriol,  copperas,  or  ferrous  sulphate,  FeSO4. 
7H2O),  sodium  sulphate  (Na2SO4,  and  Glauber's  salt, 


242  CHEMISTRY 

Na2S04.ioH2O),  and  magnesium  sulphate  (MgS04,  and 
Epsom  salts,  MgSC^.yH^O).  Sulphates  are  used  in 
many  industries. 

290.  The  test  for  sulphuric  acid  or  a  soluble  sulphate  is  the 
formation  of  the  white,  insoluble  barium  sulphate  upon  the  addi- 
tion of  barium  chloride  solution.  Both  solutions  contain  S04-ions 
which  unite  with  Ba-ions  to  form  insoluble  barium  sulphate.  The 
ionic  equation  for  the  reaction  is:  — 

Ba++  +  S04-- 


An  insoluble  sulphate  fused  on  charcoal  is  reduced  to  a  sulphide, 
which  blackens  a  moist  silver  coin,  owing  to  the  formation  of  silver 
sulphide  (Ag2S). 

291.  Fuming  Sulphuric  Acid,  H2S2O2,  is  made  by  adding  sulphur 
trioxide  to  sulphuric  acid.     This  is  the  acid  called  sulphuric  acid 
by  the  alchemists,  who  made  it  by  heating  moist  ferrous  sulphate. 
It  is  sometimes  called  Nordhausen  sulphuric  acid.      It  is  a  thick, 
brown  liquid,  which  fumes  strongly  in  the  air,  owing  to  the  escape 
of  oxides  of  sulphur.      It  is  used  in  gas  analysis  to  absorb  ethylene 
and  other  illuminants,  and  in  dyeing  to  dissolve  indigo.      If  the 
fuming  acid  is  cooled  to  9°  C.,  crystals  separate;  they  are  called 
pyrosulphuric  acid. 

292.  Sodium    ThiosulpKate,  Na2S2O3,  is  a  salt  of  an  unstable 
acid.      It  is  sometimes  incorrectly  called  sodium  hyposulphite,  or 
simply   "  hypo."      It  is  a  white,  crystalline  solid,  very  soluble  in 
water.      The  solution,  used  in  excess,  dissolves  certain  compounds 
of  silver,  i.e.  AgCl,  AgBr,   Agl,  hence  its  extensive  use  in  photo- 
graphy. 

293.  Carbon  Bisulphide    (or   Bisulphide),    CS2?    when 
pure,  is  a  clear,  colorless  liquid,  with  an  agreeable  odor, 
but  the  commercial  substance  is  yellow  and  has  an  ex- 
ceedingly offensive  odor.     It  is  poisonous.     It  is  volatile 
and  inflammable,  the  equation  for  its  combustion  being:  - 

CS2       +         3O2       =       CO2        +         2SO2 

Carbon  Bisulphide        Oxygen        Carbon  Dioxide        Sulphur  Dioxide 


CARBON  BISULPHIDE 


243 


Since  its  vapor  is  poisonous  and  inflammable,  the  liquid 
must  be  used  with  care.  This  liquid  is  practically  in- 
soluble in  water.  It  dissolves  rubber,  gums,  fats,  iodine, 
camphor,  and  some  forms  of  sulphur.  It  is  a  highly 
refracting  liquid,  and  hollow  glass  prisms  filled  with  it 
are  used  to  decompose  light.  As  a  solvent  it  is  used 
to  dissolve  pure  rubber  in  the  manufacture  of  rubber 
cement.  It  is  also  used  to  kill 
insects  on  both  living  and  dried 
plants  (e.g.  in  museums),  and  to 
exterminate  burrowing  animals, 
such  as  moles  and  woodchucks. 
Many  oils,  waxes,  and  greases 
are  extracted  by  carbon  disul- 
phide. It  is  also  used  to  manu- 
facture compounds  of  sulphur 
and  of  carbon. 


Carbon  disuJphide  is  manufactured 
by  an  electrothermal  process.  A  dia- 
gram of  the  furnace  is  shown  in  Fig. 
71.  Several  groups  of  carbon  elec- 
trodes (EE)  are  set  into  the  base  of  a 
furnace,  coke  is  packed  loosely  between 
them,  and  the  body  of  the  furnace  is 
rilled  with  charcoal  (C).  Sulphur  is  in- 


Fig. 71.  —  Electric  Furnace 
for  the  Manufacture  of 
Carbon  Disulphide. 


troduced  at  suitable  points  (S,  S,  S),  and  coke  is  fed  in  through  K,  K. 
When  the  current  passes,  the  heat  caused  by  the  resistance  offered 
by  the  coke  melts  the  sulphur,  which  unites  with  the  heated  carbon 
above  the  electrodes.  The  vapors  of  the  carbon  disulphide  escape 
through  the  pipe  (P)  and  are  condensed  in  a  special  apparatus. 


EXERCISES 

1.  Where  is  free  sulphur  found?    Name  five  or  more  native  compounds 
of  sulphur.    What  animal  and  vegetable  substances  contain  sulphur? 

2.  Give  the  name  and  formula  of  (a)  five  sulphides  and  (b)  five  sulphates. 


244  CHEMISTRY 

3.  Describe  the  lead  chamber  process  of  manufacturing  sulphuric  acid. 

4.  Starting  with  sulphur,  how  would  you  prepare  successively  sulphur 
dioxide,  sulphur  trioxide,  sulphuric  acid,  hydrogen  sulphide,  and  sulphur? 

5.  Describe  the  contact  process  of  manufacturing  sulphuric  acid. 

6.  Enumerate  the  important  uses  of  sulphuric  acid. 

7.  What  is  (a)   gypsum,  (b)  white  vitriol,  (c)   green   vitriol,  (d)  blue 
vitriol,  (e)  Glauber's  salt,  (/)  oil  of  vitriol  ? 

8.  State  the  test  for  (a)  sulphuric  acid,  (b]  sulphurous  acid,  (c)  a  soluble 
sulphate,  (d)  an  insoluble  sulphate,  (e)  a  sulphite,  (/)  hydrogen  sulphide. 

9.  Essay  topics:  (a)  Sicilian  sulphur  industry,     (b)  Louisiana  sulphur 
industry,    (c)  Uses  of  sulphur,    (d)  Uses  of  sulphuric  acid. 

10.  Explain  the  following:    (a)  Bottles  of  concentrated  sulphuric  acid 
sometimes  have  a  white  deposit  inside  the  bottle  and  a  black  one  outside. 
(b)  Bottles  of  hydrogen  sulphide  water  upon  standing  lose  the  bad  odor, 
become  cloudy,  and  have  a  white  layer  on  the  bottom,    (c)  Dishes  of  con- 
centrated sulphuric  acid  are  often  left  inside  the  case  of  a  chemical  balance. 
(d)  A  rubber  band  blackens  a  silver  coin. 

PROBLEMS 

1.  What  is  the  approximate  weight  of  sulphur  in  a  bin  200  feet  long, 
80  feet  wide,  and  90  feet  high?    (Note.  —  A  cubic  foot  of  water  weighs 
approximately  62  pounds.) 

2.  Calculate  the  weight  of  sulphur  in  (a)  a  metric  ton  of  pure  FeS2, 
(6)  a  kilogram  of  CS2,  (c)  720  gm.  of  MgSO4,  (d)  2000  Ib.  of  CaSO4. 

3.  Calculate  the  weight  of  sulphur  in  (a)  77  gm.  of  H2S,  (b)  77  1.  of  H2S 
at  o°  C.  and  760  mm.,  (c)  77  1.  of  H2S  at  20°  C.  and  755  mm. 

4.  Hydrochloric  acid  solution  having  the  specific  gravity  of  1.2  contains 
39.86  per  cent  of  HC1  (by  weight).    What  (a)  weight  and  (b)  volume  of 
hydrogen  sulphide  (at  standard  conditions)  will  be  produced  by  the  inter- 
action of  ferrous  sulphide  and  2  1.  of  this  acid? 

5.  Write  the  equation  for  the  interaction  of  sodium  sulphite  and  sul- 
phuric acid.    What  weight  and  what  volume  (standard  conditions)  of  sulphur 
dioxide  can  be  prepared  from  25  gm.  of  sodium  sulphite  (92  per  cent  pure)? 

6.  In  Problem  5  what  volume  of  sulphuric  acid  solution  is  needed,  if 
the  acid  used  has  the  specific  gravity  1.45  and  contains  55.07  per  cent  of 
H2SO4  by  weight? 

7.  An  oxide  of  sulphur  contains  50  per  cent  of  each  element.      If  the 
atomic  weights  of  oxygen  and  sulphur  are  respectively  16  and  32,  what  is 
the  molecular  formula  of  the  compound? 

8.  A  flask  filled  with  water  was  found  to  weigh  72  grams,  the  flask  alone 
weighing  22  grams.    The  same  flask  filled  with  sulphuric  acid  weighed  114 
grams.    Calculate  the  specific  gravity  of  the  sulphuric  acid. 


CHAPTER   XIX 
BORON  — BORAX  — BORIC  ACID' 

294.  Occurrence,  Preparation,  and  Properties  of  Boron. 

-  The  non-metallic  element  boron  (B)  is  never  found  free. 
Some  of  its  compounds  are  abundant,  e.g.  borax  (Na2B4O7), 
boric  acid  (H3BO3),  and  colemanite  (Ca2B6Oii.5H2O). 

Pure  boron  is  prepared  by  heating  boron  chloride  and  hydrogen 
to  a  very  high  temperature.  The  equation  is:  - 

2BiCl3  +    3H2  =     26  +      6HC1 

Boron  Chloride  Hydrogen  Boron  Hydrogen  Chloride 

Pure  boron  is  a  dark  gray  non-crystalline  solid.  It  is  very 
hard,  almost  as  hard  as  diamond.  The  melting  point  is  over 
2000°  C.  Boron  and  oxygen  unite  readily,  forming  the  stable 
compound  boric  oxide  (B2O3).  A  substance  called  "boron  sub- 
oxide"  is  obtained  by  partial  reduction  of  boric  oxide  and  is  used 
in  the  production  of  copper  castings  free  from  the  holes  that  are 
often  caused  by  oxygen  and  other  gases. 

295.  Borax,    Na^B^y,    occurs   in   large   quantities   in 
California.     Most  of  the  commercial  borax  is  made  from 
calcium   borate  (colemanite,    Ca2B6Oii.5H2O)   by  boiling 
it  with  sodium  carbonate  and  separating  the  borax  by 
crystallization.     The  equation  for  the  reaction  is:  - 

2Ca2B5On  +  4Na2CO3  +  H2O  =  3Na2B4O7  +  4CaCO3  +  2NaOH 

Calcium  Sodium  Water  Borax  Calcium  Sodium 

Borate  Carbonate  Carbonate  Hydroxide 

Borax  is  a  white  crystallized  solid,  having  five  or  ten 
molecules  of  water  of  crystallization;  the  usual  commercial 
form  is  the  powdered  crystals.  The  crystals  readily 
effloresce  and  crumble  in  the  air.  When  heated,  crystal- 


246  CHEMISTRY 

lized  borax  loses  its  water  of  crystallization  and  swells  into 
a  white  porous  mass  which  finally  melts  into  a  glass-like 
solid.  This  glassy  borax  dissolves  metallic  oxides  and  is 
colored  by  them.  If  the  borax  is  melted  on  the  end  of  a 
looped  wire,  the  transparent  globule  is  called  a  borax 
bead.  These  beads  usually  assume  different  colors  after 
being  heated,  in  an  oxidizing  or  a  reducing  flame,  and  the 
colors  are  characteristic  of  the  metals.  Thus,  a  copper 
bead  is  made  blue-green  by  an  oxidizing  flame  and  red 
by  a  reducing  flame.  (See  Part  II,  Exps.  142-144.) 

Borax  has  various  uses.  Its  chief  use  is  in  the  manu- 
facture of  enamels  for  coating  iron  ware.  The  so-called 
"  granite"  or  "  agate"  ware  and  "porcelain-lined"  vessels 
are  made  of  iron  coated  with  an  easily  fused  glass  called 
enamel.  Some  is  used  for  preserving  meat,  fish,  cheese, 
and  other  foods,  because  it  prevents  the  growth  of  certain 
bacteria  which  produce  decay.  A  solution  of  borax  has 
an  alkaline  reaction;  hence  it  is  sometimes  used  instead 
of  soap  as  a  cleansing  agent.  Some  soaps  contain  borax. 
In  laundries  borax  is  used  to  soften  hard  water  and  to 
improve  the  finish  of  ironed  goods.  Its  power  to  dis- 
solve oxides  adapts  borax  for  use  in  soldering  and  welding 
metals.  Solder  adheres  only  to  clean  metals,  so  a  little 
borax  is  used  to  dissolve  the  film  of  oxide  on  the  surfaces 
to  be  joined.  Considerable  is  consumed  as  an  ingredient 
of  ointments,  lotions,  and  toilet  powders. 

296.  Boric  Acid  (or  boracic  acid),  H3BO3,  is  contained  in  the 
waters  and  steam  of  certain  hot  springs,  especially  in  Italy.  The 
acid  is  also  prepared  by  the  interaction  of  sulphuric  acid  and  borax. 
The  equation  for  the  reaction  is:  — 


H2SO4     +  Na2B407     +     sH2O   =  ^BOs    +      Na2SO4 

Sulphuric  Borax  Water  Boric  Sodium 

Acid  Acid  Sulphate 


BORIC  ACID  247 

Boric  acid  crystallizes  in  lustrous,  white  flakes,  which  feel 
greasy.  It  dissolves  slightly  in  cold  water,  readily  in  hot  water, 
and  in  alcohol.  Boric  acid  responds  only  feebly  to  the  usual  tests 
for  acids;  it  forms  no  salts. 

Boric  acid  is  used  to  some  extent  in  the  manufacture  of  enamels 
and  glazes  for  pottery.  Solutions  are  used  as  an  antiseptic  wash 
for  the  eyes  and  for  parts  of  the  body  affected  by  blood  poisoning; 
the  powder  is  sometimes  used  to  relieve  burns.  Formerly  con- 
siderable was  used  for  preserving  meat,  fish,  milk,  butter,  beer, 
and  wine. 

When  an  alcoholic  solution  of  boric  acid  is  burned,  a  boron 
compound  colors  the  vapor  green.  Since  many  boron  compounds 
can  be  converted  into  boric  acid,  the  green  flame  serves  as  the 
test  for  boron  compounds. 

EXERCISES 

1.  Give  the  name  and  formula  of  three  native  compounds  of  boron. 

2.  How  is  borax  manufactured?    State  its  properties  and  uses. 

3.  Describe  a  borax  bead.    State  and  illustrate  its  use. 

4.  Essay  topics:   (a)  Boric  acid  industry  in  Tuscany,    (b)  Uses  of  borax. 
(c)  Manufacture  of  enamel  ware,    (d)  Borax  industry  in  the  United  States. 

PROBLEMS 

1..  How  many  grams  of  boron  can  be  made  by  reducing  (a)  250  gm.  of 
boron  trichloride  with  hydrogen  and  (b)  200  gm.  of  boron  trioxide  with 
magnesium? 

2.  Calculate  the  per  cent  of  boron  in  (a)  colemanite,  (b)  boron  trichloride, 
(c)  borax,  (d)  boric  acid,  (e)  sodium  metaborate  (NaBO2),  (/)  boron  tri- 
oxide. 

3.  HOW  many  grams  of  boric  acid  can  be  made  from  60  gm.  of  borax 
(95  per  cent  pure)? 

4.  The  formula  of  tetraboric  acid  is  H2B4O7  and  of  metaboric  acid  is 
HBOa.    Write  the  formulas  of  the  salts  of  each  acid  corresponding  to  potas- 
sium, copper  (ic  and  ous), barium,  aluminium, magnesium,  cobalt,  chromium. 
Calculate  the  per  cent  of  boron  in  any  three  of  these  salts. 


CHAPTER   XX 

SILICON  — SILICON  DIOXIDE  —  SILICIC  ACID   AND 
SILICATES  — GLASS 

297.  Occurrence. -- The  element  silicon  (Si)  does  not 
occur  free  in  nature,  but  some  of  its  compounds  are  very 
abundant,  especially  silicon  dioxide  (SiO2)  and  silicates. 
Sand  is  silicon  dioxide  and  many  rocks  are  made  up  of 
silicates.      Silicon  comes  next  to  oxygen  in  abundance, 
about  one  fourth  of  the  earth's  crust  being  silicon. 

298.  Preparation.  —  Silicon  is  prepared  on  a  commer- 
cial scale  by  heating  a  mixture  of  silicon  dioxide  (i.e.  sand) 
and  carbon  (in  the  form  of  coke)  in  an  electric  furnace. 
The  equation  for  the  reaction  is:  - 

SiO2       +       2C     =    Si     +  2  CO 

Silicon  Dioxide         Carbon         Silicon         Carbon  Monoxide 

299.  Properties. -- Thus  prepared,  silicon  is  a  brittle, 
gray-black,  lustrous,  metallic  looking  solid.     It  melts  at 
about  1400  °  C.     The  specific  gravity  is  about  2.37.     It 
is  hard  enough  to  produce  scratches  on  glass,  being  almost 
as  hard  as  quartz. 

Silicon  at  high  temperatures  combines  with  oxygen, 
though  not  readily,  to  form  silicon  dioxide  (SiO2);  with 
other  elements  it  forms  compounds  often  called  silicides, 
e.g.  carbon  silicide  (CSi).  Chlorine  and  elements  related 
to  it,  especially  fluorine,  combine  with  silicon  to  form 
volatile  compounds.  Sodium  hydroxide  and  potassium 
hydroxide  interact  with  silicon,  thus:  - 


SILICON  249 

4NaOH     +      Si  Na4SiO4     +      2H2 

Sodium  Hydroxide       Silicon        Sodium  Silicate        Hydrogen 

Water  containing  alkaline  substances,  such  as  sodium 
carbonate,  acts  slowly  upon  silicon  in  much  the  same  way. 
Most  acids  produce  little  or  no  effect  upon  silicon,  hydro- 
fluoric acid  being  the  main  exception.  (See  Part  II,  Exp. 
146.) 

300.  Uses.  —  Silicon  is  made  into  vessels  and  parts  of 
apparatus  designed  to  withstand  the  action  of  acids  and 
other  liquids.    Some  is  used  to  improve  castings  by  remov- 
ing the  gases  that  cause  small  holes. 

301.  Silicon  Dioxide  or  Silica,  SiO2,  is  the  most  com- 
mon and  important  compound  of  silicon.     Sand,  gravel, 
sandstone,  quartzite,  and  the  numerous  varieties  of  quartz 
are    almost    wholly    silicon    dioxide;     many    rocks,    e.g. 
granite  and  gneiss,  contain  silica  as  an  essential  ingredient. 
Opal  is  a  variety  of  silica  containing  varying  per  cents  of 
combined  water.     Petrified  or  silicified  wood  is  the  rem- 
nant of  wood  in  which  the  fiber  has  been  replaced  by 
silica.      There    is    a 

petrified  forest  in 

Arizona.     Infusorial 

or    diatomaceous 

earth  is  largely  silica 

and  consists  of  the 

shells  of  minute  organisms  called  diatoms,  many  being 

of  delicate  and  beautiful  structure  (Fig.  72). 

302.  Quartz  is  the  commonest  form  of  silicon  dioxide. 
Pure  quartz  is  colorless  and  transparent,  and  is  called 
rock  crystal.     It  is  crystalline  and  is  frequently  found  as 
single  crystals  which  consist  usually  of  a  six-sided  prism 
with  a  six-sided  pyramid  at  one  or  both  ends;  the  crystals 


250  CHEMISTRY 

are  sometimes  distorted  or  complex  (Fig.  73).  There 
are  many  varieties  of  quartz  which  differ  in  color  and 
structure.  Among  the  crystalline  vari- 
eties are  the  clear,  colorless  rock  crystal, 
the  purple  amethyst,  and  the  rose,  yel- 
low, glassy,  milky,  and  smoky  forms. 
Varieties  imperfectly  crystalline  are  the 
Fig.  73.  —  Quartz  waxlike  chalcedony,  the  various  forms  of 
agate  having  different  colored  layers,  the 
reddish  brown  carnelian,  the  black  and  white  onyx,  the 
red  or  brown  jasper,  the  brown  or  black  flint. 

303.  Properties  of  Silicon  Dioxide.  —  Quartz  and  most 
crystalline  varieties  of  silica  are  harder  than  other  common 
substances.  The  specific  gravity  of  quartz  is  2.66.  Quartz 
and  other  varieties  of  silica  are  infusible,  except  in  the 
oxyhydrogen  flame  and  the  electric  furnace.  If  pure  silica 
is  fused  with  certain  precautions,  the  viscous  mass  can  be 
drawn  into  elastic  threads,  which  are  used  to  suspend 
delicate  parts  of  electrical  instruments;  it  can  also  be 
shaped  into  tubes,  flasks,  crucibles,  etc.,  which  do  not 
crack  when  heated  to  a  high  temperature  and  then  sud- 
denly cooled. 

The  different  forms  of  silicon  dioxide  do  not  dissolve  in 
water.  They  resist  the  action  of  acids,  except  hydrofluoric 
acid,  which  transforms  them  into  a  volatile  fluoride  (silicon 
tetrafluoride,  SiF4) .  They  are  converted  into  soluble  sili- 
cates when  heated  in  water  containing  alkaline  substances 
or  when  fused  with  the  hydroxides  or  carbonates  of  sodium 
and  potassium.  Thus,  wjien  fine  sand  is  fused  with 
sodium  carbonate  the  equation  for  the  reaction  is:  - 

Si02     +    Na2CO3    =    Na2SiO3     +        CO2 

Silicon  Sodium  Sodium  Carbon 

Dioxide  Carbonate  Silicate  Dioxide 


SILICON  251 

304.  Uses  of  Silicon  Dioxide.  —  Sandstone  and  quartz- 
ite  are  used  as  building  stones,    and  some  varieties  of 
sandstone  are  shaped  into  grindstones  and  whetstones. 
Immense  quantities  of  sand  are  used  in  making  sandpaper, 
glass,  porcelain,  cement,  and  mortar.    Different  grades  of 
sand  are  used  as  grinding  and  polishing -material.     Glass 
is  polished  by  rubbing  it  with  fine  wet  sand;    it  is  also 
roughened  and  cut  by  blowing  or  " blasting"  fine  sand 
against  it.     Glass  stoppers  for  bottles  used  in  the  labora- 
tory are  "ground"  with  sand.     Many  of  the  varieties  of 
quartz  are  cut  and  polished  into  ornaments  and  gems. 
Infusorial  earth  is  used  as  a  polishing  powder  for  metals 
(" electro-silicon"    being   the   commercial    name    of   one 
kind)    and   in   making   scouring   soaps,    cement,    soluble 
glass,   and  refractory  brick;    considerable  is  used  as  a 
packing  material  for  steam  pipes  and  boilers  and,  owing 
to  the  hollow  structure  of  its  minute  particles,   as  an 
absorbent  of  nitroglycerin  in  the  manufacture  of  dyna- 
mite.    Fused  silica  vessels  and  other  apparatus  are  used 
in  industries  which  require  material  resistant  to  chemicals. 

305.  Silica  and  Life.  —  The  ash  of  many  plants  contains  silica. 
The  ash  of  rye  and  wheat  straws  and  of  potato  stems  contains  from 
40  to  70  per  cent  of  silica.     Silica  is  also  an  ingredient  of  the  ash  of  the 
feathers  and  hair  of  animals.     Certain  minute  aquatic  organisms 
have  a  shell  or  coating  of  silica  (see  Diatoms,  above).     Probably  in 
plants  and  animals  the  silicon  is  present  as  a  constituent  of  complex 
compounds,  and  when  these  are  burned  the  silicon  finally  forms  the 
oxide. 

306.  Silicic   Acid   and   Silicates.  —  When  fine   silicon 
dioxide  is  fused  with  sodium  carbonate,  sodium  silicate 
is  formed,  as  already  stated.    Sodium  silicate,  unlike  most 
silicates,  dissolves  in  water,  and  when  hydrochloric  acid 
is  added  to  concentrated  sodium  silicate  solution,  a  white 


252  CHEMISTRY 

jelly-like  precipitate  called  silicic  acid  is  formed.  The 
equation  usually  given  for  the  reaction  is-:  - 

Na2Si03     +     2HC1     +     H2SiO3     +     2NaCl 

Sodium  Hydrochloric  Silicic  Sodium 

Silicate  Acid  Acid  Chloride 

Probably  other  salts  of  silicon  are  formed  by  fusing  silicon 
dioxide  with  such  substances  as  the  hydroxides  and  car- 
bonates of  sodium  and  potassium,  and  when  the  solution 
is  treated  with  acids  like  hydrochloric  acid  the  precipitate 
is  a  mixture  of  meta-  and  orthosilicic  acids  (H2SiO3  and 
H4SiO4) .  There  are  several  silicic  acids,  though  they  have 
not  been  prepared  in  the  pure  state.  Their  salts,  however, 
are  well  known;  indeed,  the  silicates  are  among  the  com- 
monest substances  in  the  earth's  crust.  Many  minerals 
and  rocks  consist  wholly  or  largely  of  silicates  of  the  metals 
aluminium,  iron,  calcium,  potassium,  sodium,  and  mag- 
nesium. Some  common  silicates  are  feldspar,  mica,  horn- 
blende, clay,  beryl,  garnet,  serpentine,  and  talc;  the  lava 
which  is  ejected  by  volcanoes  consists  largely  of  fused 
silicates  from  the  interior  of  the  earth. 

The  simplest  silicic  acids  are  metasilicic  acid  (H2SiO3) 
and  orthosilicic  acid  (H4SiO4).  .The  ortho-acid  passes 
into  the  meta-acid  upon  drying,  thus :  - 

H4Si04       =      H2Si03       +  H20 

Orthosilicic  Acid        Metasilicic  Acid        Water 

And  the  metasilicic  acid  when  heated  decomposes  into 
silicon  dioxide  and  water,  thus :  - 

H2SiO3  SiO2        +  H2O 

Metasilicic  Acid        Silicon  Dioxide       Water 

These  two  silicic  acids  are  closely  related  to  each  other 
and  to  silicon  dioxide,  which  may  be  regarded  as  their 
anhydride,  though  the  acid  cannot  be  formed  by  dis- 


SILICON  253 

solving  silica  in  water.  Other  silicic  acids  are  complex, 
and,  although  they  have  not  been  isolated,  they  may  be 
conveniently  thought  of  as  one  or  more  molecules  of 
orthosilicic  acid  minus  one  or  more  molecules  of  water. 
Thus,  we  might  have  the  following  hypothetical  silicic 
acids :  — 

2H4SiO4  -     H2O  =  H6Si2O7,  or  Disilicic  Acid 
3H4SiO4  -  4H2O  =  H4Si3O8,  or  Trisilicic  Acid 

Many  of  the  silicates  in  the  earth  are  salts  of  these  poly- 
silicic  acids,  as  they  are  often  called.  For  example,  ser- 
pentine (Mg3Si2O8)  is  a  salt  of  disilicic  acid  (H6Si2O7),  and 
orthoclase  (KAlSi3O8)  of  trisilicic  acid  (H4Si3O8).  Illus- 
trations of  the  simpler  silicates  are  the  metasilicate  wol- 
lastonite  (CaSi03)  and  the  orthosilicate  zircon  (ZrSiO4). 

Sodium  and  potassium  silicates  are  the  only  salts  readily 
soluble  in  water,  and  the  thick,  sirupy  solution  is  often 
called  water  glass  or  soluble  glass,  because  it  hardens  to 
a  glassy  solid  when  exposed  to  air.  The  solution  is  used 
in  making  yellow  soaps,  cements,  and  artificial  stone,  to 
fix  colors  in  frescoing  and  calico  printing,  and  to  render 
cloth,  wood,  and  paper  fireproof.  Considerable  is  used 
to  preserve  eggs ;  the  eggs  are  put  in  a  vessel  and  covered 
with  the  solution,  which  keeps  the  air  from  penetrating 
the  shell  and  thus  prevents  decomposition.  (See  Part  II, 
Exp.  148.) 

Naturally  occurring  silicates  are  called  insoluble; 
nevertheless  they  dissolve  very  slowly,  especially  through 
the  joint  action  of  carbon  dioxide  and  water.  Even  the 
moist  carbon  dioxide  in  the  air  slowly  decomposes  silicates, 
and  as  a  result  of  this  gradual  but  very  extensive  chemical 
change,  hard  rocks  like  granite  disintegrate,  the  more 
soluble  products  being  washed  away  and  the  less  soluble 


254  CHEMISTRY 

ones  being  left  behind  as  clay,  sand,  etc.  This  transforma- 
tion, which  is  called  weathering,  is  assisted  by  the  action 
of  water,  especially  alternate  freezing  and  thawing  as  well 
as  dissolving. 

The  decomposition  of  silicates  by  the  joint  action  of 
carbon  dioxide  and  water  (i.e.  by  carbonic  acid  (H2CO3)) 
is  due  to  the  fact  that  silicic  acid  is  a  weak  acid.  This 
property  is  strikingly  shown  by  the  formation  of  basins 
and  terraces  of  silica  in  Yellowstone  Park.  The  hot  waters 
of  some  of  the  springs  in  the  Park  contain  dissolved  alka- 
line silicates  (e.g.  Na4Si04)  which  are  transformed  to 
silicic  acid  by  the  carbon  dioxide  of  the  air;  the  silicic 
acid  which  is  left  where  the  waters  overflow  gradually 
loses  water  and  builds,  up  a  deposit  of  silica,  often  very 
beautiful.  Silica  thus  deposited  is  called  geyserite  or 
siliceous  sinter. 

307.  Colloidal  Silicic  Acid.  —  Sodium  silicate  and  hydrochloric 
acid  do  not  always  interact  as  described  above  (306).  If  the  sodium 
silicate  solution  is  dilute,  or  the  hydrochloric  acid  is  concentrated  or 
in  excess,  or  the  mixing  is  done  quickly,  then  the  silicic  acid  which 
is  formed  remains  in  solution.  This  solution,  however,  is  not  a  solu- 
tion of  the  gelatinous  silicic  acid  described  above,  for  the  latter  is 
insoluble  in  water.  Moreover,  if  this  solution  is  evaporated,  the 
silicic  acid  separates  as  a  gelatinous  mass  which  will  not  redissolve 
upon  addition  of  water.  The  silicic  acid  in  the  solution  is  in  the 
colloidal  state  or  in  colloidal  solution.  It  cannot  be  filtered  out  by 
passing  the  liquid  through  paper  in  the  usual  way. 

Many  substances,  e.g.  glue,  starch,  gelatin,  albumin,  form  colloidal 
solutions.  In  our  bodies  colloidal  solutions  play  an  important  part 
in  vital  processes.  Colloidal  solutions  in  all  probability  are  really  not 
solutions  at  all,  but  liquids  in  which  the  colloid  is  suspended  as  very 
fine  particles,  too  fine  to  be  retained  by  filter  paper.  This  view  is 
substantiated  by  the  fact  that  a  beam  of  light  when  passed  through 
a  colloidal  solution  is  made  bright  by  the  particles  just  as  it  is  by  the 
dust  in  a  room. 


GLASS  255 

308.  Other  Silicon  Compounds.  —  When  hydrofluoric  acid  in- 
teracts with  .silicon  dioxide  or  silicates,  silicon  tetrafluoride  (SiF4) 
is  formed.  Thus,  with  silicon  dioxide  the  equation  is:  — 

SiO2     +    4HF     =     SiF4     +     2H2O 

Silicon        Hydrofluoric  Silicon  Water 

Dioxide  Acid  Tetrafluoride 

Silicon  tetrafluoride  is  a  colorless  gas  which  fumes  in  moist  air  owing 
to  interaction  with  water.  The  equation  for  the  reaction  is:  — 

3SiF4     +    3H20     =     H2Si03     +     2H2SiF6 

Silicon  Water  Silicic  Hydrofluosilicic 

Tetrafluoride  Acid  Acid 

The  hydrofluosilicic  acid  (sometimes  called  simply  fluosilicic  acid) 
remains  in  solution,  while  the  silicic  acid  is  precipitated.  The  forma- 
tion of  the  white  gelatinous  silicic  acid  when  the  gases  from  the  inter- 
action of  hydrofluoric  acid  and  a  compound  of  silicon  are  led  into 
water  is  often  used  as  a  test  for  silicon.  (See  Part  II,  Exps.  147, 150.) 
For  carbon  silicide  (or  silicon  carbide,  carborundum,  CSi)  see  200. 


Glass 

309.  Glass    and    its    Ingredients.  —  Glass    is    a    non- 
crystalline,  rigid  substance,  which  consists  essentially  of 
a  mixture  of  certain  silicates  and  silica.     It  results  from 
the  cooling  of  a  fused  mixture  of  sand  and  several  other 
substances. 

Besides  sand,  the  substances  needed  in  making  glass 
are  a  calcium  or  a  lead  compound  and  a  sodium  or  a  potas- 
sium compound.  Besides  these  fundamental  ingredients, 
other  substances  are  used,  e.g.  sodium  nitrate,  carbon, 
arsenic  trioxide,  and  manganese  dioxide.  The  general 
proportions  of  the  ingredients  used  in  making  several 
types  of  American  glass  are  shown  in  the  Table  of  Pro- 
portions of  Ingredients  for  Glass  on  page  256. 

310.  Manufacture   of   Glass.  —  The    process   of    glass 
making  consists  in  melting  a  mixture  of  the  proper  in- 


256  CHEMISTRY 

TABLE  OF  PROPORTIONS  OF  INGREDIENTS  FOR  GLASS 


Ingredient 

Plate 

glass 

Win- 
dow 
glass 

Green 
bottle 
glass 

Lead 
glass 

Sand  (SiO2)   

IOO 

IOO 

IOO 

Sodium  sulphate  (Na2SC>4)    
Sodium  carbonate  (Na2COs) 

36 

42 

38 

Limestone  (CaCO3)    

24 

AO 

•2  A 

Carbon  (C)    

7r 

6 

Arsenic  trioxide  (AssOs)  

I 

2 

T  r 

Potassium  carbonate  (K2CO3)    .... 
Red  Lead  (Pb3O4)  

34 
d8 

Sodium  nitrate  (NaNOs) 

6 

Manganese  dioxide  (MnO2)    

06 

Antimony  (Sb)    . 

gradients  in  a  fire-clay  pot.  There  are  four  main  kinds 
of  glass  and  many  varieties  of  each.  In  the  order  of 
their  fusibility  and  beginning  with  the  softest,  the  four 
kinds  are:  (i)  Sodium-lead  glass,  (2)  potassium-lead 
glass,  (3)  sodium-calcium  glass,  and  (4)  potassium-cal- 
cium glass.  Glass  made  from  a  lead  compound  is  often 
called  flint  glass;  it  is  lustrous,  refracts  light  to  a  high 
degree,  and  is  made  into  lenses  for  optical  instruments 
and  into  shades  for  electric  and  gas  lights.  Cut  glass  is 
a  lead  glass.  Window,  plate,  crown,  table,  and  bottle 
glass  are  varieties  of  sodium-calcium  glass.  This  kind  of 
glass  softens  when  heated  and  the  flame  becomes  yellow 
from  the  sodium,  hence  it  is  often  called  soft  glass  or 
soda  glass.  Glass  tubing  and  much  chemical  glassware 
is  sodium-calcium  glass.  Bohemian  or  hard  glass  is  a 
potassium-calcium  glass,  and  is  used  in  making  chemical 
apparatus  designed  to  withstand  great  heat.  Soft  glass 
is  slightly  soluble  in  water,  but  hard  glass  is  less  so,  hence 


GLASS  257 

special  varieties  of  hard  glass  are  often  made  into  ap- 
paratus which  resists  the  solvent  action  of  water  and 
chemical  reagents.  Jena  glass  is  one  variety  of  hard  glass. 
(See  Part  II,  Exp.  151.) 

In  making  glass  objects  the  molten  glass  is  manipulated  in  various 
ways.  Bottles,  for  example,  are  made  by  gathering  a  mass  of  the 
plastic  glass  on  the  end  of  a  long  tube,  called  a  glass  blowpipe,  blowing 
the  glass  into  a  preliminary  shape,  then  lowering  it  into  a  mold  and 
blowing  until  the  glass  fills  the  mold,  and  finally  opening  the  mold 
and  lifting  out  the  bottle;  by  subsequent  operations  the  neck  of  the 
bottle  is  shaped  and  finished.  Lamp  chimneys  are  made  by  blowing 
the  lump  of  glass  into  the  desired  shape  and  simultaneously  swinging 
and  twisting  the  plastic  mass;  no  mold  is  used  and  considerable  skill 
is  necessary  to  produce  the  proper  shape.  Window  glass  is  made  by 
blowing  a  lump  of  glass  into  a  hollow  globe  and  then  into  a  cylinder; 
this,  on  being  opened  at  both  ends  and  cut  lengthwise,  spreads  out  flat. 
Plate  glass  is  made  by  pouring  the  molten  glass  upon  a  large  table, 
rolling  it  with  a  hot  iron  roller,  and  subsequently  grinding  and  polish- 
ing it  until  the  surfaces  are  parallel.  Some  cheap  articles  are  made 
by  pressing  plastic  glass  with  a  die  and  certain  kinds  of  inexpensive 
hollow  ware  are  blown  into  shape  by  machinery. 

Glass  must  be  cooled  slowly  to  prevent  brittleness.  This  opera- 
tion is  called  annealing,  and  is  accomplished  by  passing  the  objects 
slowly  through  a  furnace  in  which  the  temperature  is  gradually 
lowered. 

Glass  is  colored  by  adding  different  substances  to  the  molten  mass. 
Iron,  chromium,  and  certain  copper  compounds  make  it  green,  the 
green  color  of  many  bottles  and  fruit  jars  being  due  to  iron  com- 
pounds in  the  original  materials;  selenium  produces  a  red  color,  and 
selenium  glass  is  now  used  in  red  signal  lanterns;  white  glass  is  made 
by  adding  fluor  spar,  cryolite,  or  calcium  phosphate;  stained  glass  is 
ordinary  glass  to  which  fusible  pigments  are  applied  with  a  brush 
and  then  fixed  by  heat. 


258  CHEMISTRY 

EXERCISES 

1.  In  what  compounds  is  silicon  found  in  nature?    What  proportion  of 
the  earth's  crust  is  combined  silicon?     Compare  the  abundance  of  silicon 
with  that  of  other  elements. 

2.  Describe  the  different  varieties  of  quartz.    Summarize  the  properties 
of  quartz.     How  can  it  be  readily  distinguished  from  other  minerals  and 
from  rocks? 

3.  Discuss  the  cycle  of  silicon  dioxide  in  nature. 

4.  Name  several  minerals  and  rocks  which  are  silicates. 

6.   Describe  the  formation,  state  the  uses,  and  enumerate  the  properties 
of  water  glass. 

6.  Starting  with  silicon,  how  would  you   prepare  successively  silicon 
dioxide,  sodium  silicate,  silicic  acid,  silicon  dioxide,  silicon? 

7.  Essay  topics:    (a)  Diatoms,  their   formation,  deposition,   and    uses. 
(b)  Manufacture  of  glass,    (c)  Quartz,    (d)  Silicon  dioxide  as  an  abrasive. 

PROBLEMS 

1.  Calculate  the  per  cent  of  silicon  in  (a)  orthosilicic  acid,  (b)  metasilicic 
acid,  (c)  potassium  feldspar,  KAlSi3O8,  (d)  sodium  feldspar,  NaAlSi3O8. 

2.  (a)  How  much  sodium  silicate  can  be  made  from  a  metric  ton  of  sand 
(85  per  cent  pure)?     (b)  How  much  potassium  silicate? 

3.  Write  the  formula  of   (a)  silicon  chloride,   (b)  silicon  bromide,  (c) 
cupric  silicide,    (d)  cuprous  silicide,  (e)  sodium  metasilicate,  (/)  sodium 
orthosilicate,    (g)  magnesium    metasilicate,    (h)   magnesium    orthosilicate, 
(t)   aluminium  metasilicate,  (j)  ferrous  orthosilicate.     Calculate   the  per 
cent  of  silicon  in  any  three  of  these  compounds. 

4.  Scheele  found  that  .6738  gm.  of  silicon  tetrachloride  gave  2.277  gm- 
of  silver  chloride.     Calculate  the  atomic  weight  of  silicon.     (Equation  is 
SiCl4  +  4AgN03  +  2H20  =  SiO2  +  4AgCl  +  4HNO3.) 


CHAPTER  XXI 

CLASSIFICATION   OF    THE   ELEMENTS  —  METALS   AND 
NON-METALS  —  PERIODIC   CLASSIFICATION 

311.  Classification  of  the  Elements.  —  In  the  preced- 
ing chapters  emphasis  was  laid  on  individual  elements, 
e.g.   oxygen,   hydrogen,   chlorine,   nitrogen,   carbon,   and 
sulphur;   others  were  mentioned,  e.g.  sodium,  potassium, 
calcium.     Certain  ones  are  quite  similar  and  can  be  put 
in  classes  which  reveal  many  fundamental  relations  of 
the  elements  and  also  help  us  in  studying  chemistry. 

Metals  and  Non-Metals 

312.  Metals   and    Non-Metals. — About   the    time    of 
Lavoisier    (1743-1794)    the   elements   were   divided  into 
metals    and    non-metals.     Those    elements    were    called 
metals  which  were  hard,,  lustrous,  heavy,  and  good  con- 
ductors of  heat,  while  the  others  were  called  non-metals. 
This  classification  has  been  retained,  and  as  additional 
elements  were  discovered  they  have  been  placed  in  the 
proper  division.     (See  Table,  page  260.) 

This  classification  is  not  accurate,  since  certain  elements 
act  as  metals  under  some  conditions  and  as  non-metals 
under  other  conditions.  These  border-line  elements  are 
sometimes  called  metalloids;  they  are  aluminium,  tin, 
antimony,  arsenic,  chromium,  and  manganese. 


260  CHEMISTRY 

TABLE  OF  IMPORTANT  METALS  AND  NON-METALS 


Metals 

Non-Metals 

Sodium 

Magnesium 

Hydrogen 

Oxygen 

Potassium 

Zinc 

Chromium 

— 

Sulphur 

— 

Cadmium 

— 

Boron 

— 

Copper 

Mercury 

Manganese 

— 

Fluorine 

Silver 

— 

— 

Carbon 

Chlorine 

Gold 

Aluminium 

Iron 

Silicon 

Bromine 

— 

— 

Cobalt 

— 

Iodine 

Calcium 

Tin 

Nickel 

Nitrogen 

Strontium 

Lead 

— 

Phosphorus 

Barium 

Antimony 

Platinum 

Arsenic 

Bismuth 

Periodic  Classification 

313.  Periodic  Classification.  —  In  1869  the  Russian 
chemist  Mendelejeff  published  the  periodic  classification 
of  the  elements.  It  is  based  on  a  relation  between  the 
properties  of  all  the  elements  and  their  atomic  weights. 
The  scheme  of  the  classification  is  substantially  as  fol- 
lows: If  the  elements  beginning  with  helium  are  arranged 
in  the  order  of  their  increasing  atomic  weights,  a  series 
results  in  which  similar  or  closely  related  elements  occur 
at  regular  intervals.  That  is,  the  series  breaks  up  natu- 
rally into  several  periods,  and  hence  the  system  of  classi- 
fication is  called  periodic.  If  the  series  is  divided  into 
these  periods  and  the  periods  are  placed  below  each  other, 
a  table  is  obtained  in  which  the  elements  can  be  viewed 
in  three  ways:  first  as  a  long  consecutive  series,  second 
as  groups,  and  third  as  periods.  The  groups  are  in  the 
vertical  columns  and  are  designated  by  the  numerals  O  to 
VIII;  the  groups  (except  Group  O)  are  subdivided  into 
families.  The  periods  are  in  horizontal  rows  and  are  num- 
bered from  i  to  12.  Such  a  table  is  shown  on  page  261. 


CLASSIFICATION   OF    THE    ELEMENTS        261 


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262  CHEMISTRY 

314.  Groups  and  Families. — An  examination  of  the 
elements  in  the  groups  shown  in  the  vertical  columns  of 
the  periodic  table  reveals  many  interesting  facts.  First, 
elements  which  resemble  one  another  are  found  in  the 
same  group.  Second,  in  a  given  group  certain  elements 
are  more  closely  related  than  others,  giving  rise  to 
sub-groups  or  families.  In  some  of  these  families  the 
similarity  of  the  members  is  very  marked,  e.g.  the 
sodium  family  (Li,  Na,  K,  etc.)  and  the  halogen  family 
(F,  Cl,  Br,  I).  We  shall  find,  as  we  proceed,  that  in 
the  groups,  and  especially  in  the  families,  the  proper- 
ties of  the  elements  show  a  gradation  in  the  same  order 
as  their  atomic  weights  (see  331). 

For  convenience,  the  elements  discussed  in  this  book 
have  been  arranged  here  as  groups  and  families. 

Group  0.  Inert- elements  or  argon  family — helium,  neon,  argon, 
krypton,  xenon. 

Group  I.  Alkali  metals  or  sodium  family  —  lithium,  sodium, 
potassium. 

Univalent   heavy    metals   or  copper   family  —  copper,    silver, 
gold. 

Group  II.  Alkaline  earth  metals  or  calcium  family  —  calcium, 
strontium,  barium,  radium. 

Bivalent  heavy  metals  or  zinc  family  —  magnesium,  zinc,  cad- 
mium, mercury. 

Group  III.     Boron  family  —  boron. 

Earth  metals  or  aluminium  family  —  aluminium. 
Group  IV.     Quadrivalent  non-metals  or  carbon  family  —  carbon, 
silicon. 

Quadrivalent  metals  or  tin  family  —  tin,  lead. 
Group  V.     Quinquivalent  non-metals  and  metals  or  nitrogen  family 
• — nitrogen,  phosphorus,  arsenic,  antimony,  bismuth. 

Group  VI.  Hexavalent  metals  or  chromium  family  —  chromium, 
tungsten,  uranium. 

Hexavalent  non-metals  or  oxygen  family  —  oxygen,  sulphur. 
Group  VII.     Manganese  family  —  manganese. 


CLASSIFICATION  OF   THE   ELEMENTS       263 

Halogen  elements  or  chlorine  family  —  fluorine,  chlorine,  bro- 
mine, iodine. 

Group  VIII.     Iron  family  —  iron,  cobalt,  nickel. 
Platinum  family  —  platinum. 

This  arrangement  should  be  studied  carefully  and 
referred  to  constantly  as  the  different  elements  are 
studied.  Comparison  with  the  periodic  table  and  with 
the  table  of  metals  and  non-metals  shows  that  related 
elements  have  been  grouped  in  the  same  way. 

315.  Periods  and  the  Periodic  Law.  —  The  elements  in 
the  same  horizontal  row  in  the  periodic  table  belong  to 
the  same  period.  The  periodic  variations  of  the  prop- 
erties of  certain  typical  elements  may  be  illustrated  by 
periods  2  and  3.  Ignoring  the  argon  group  (Group  O), 
which  is  somewhat  anomalous,  and  beginning  with 
lithium,  the  metallic  properties  found  in  lithium  are 
weaker  in  beryllium  and  still  weaker  in  boron;  while  the 
feeble  non-metallic  properties  of  carbon  become  more 
marked  as  we  pass  on  through  nitrogen  and  oxygen  until 
they  reach  a  maximum  in  fluorine.  The  next  element  is 
sodium,  in  which  the  metallic  properties  reappear;  its 
place  therefore  is  beneath  lithium.  Proceeding  onward 
from  sodium,  the  same  decrease  of  basic  and  increase  of 
acid  properties  is  noticed  until  potassium  is  reached,  and 
here  again  the  marked  metallic  element  reappears;  hence 
sodium  comes  under  lithium,  magnesium  under  beryllium, 
aluminium  under  boron,  carbon  under  silicon,  and  so  on 
through  chlorine  to  potassium,  which  comes  under  its 
closely  related  element  sodium.  There  is  no  sudden 
change  in  properties  until  we  pass  from  one  period  to 
the  next.  Thus,  fluorine  at  the  end  of  period  '2  forms  a 
strong  acid,  but  sodium  at  the  beginning  of  period  3 
forms  a  strong  base.  Similarly,  chlorine  is  strongly 


264  CHEMISTRY 

acidic,  but  potassium,  which  is  the  first  metal  in  the  next 
period,  is  markedly  basic;  chlorine  is  a  typical  non- 
metal,  while  potassium  is  a  typical  metal.  Not  all  the 
periods  are  as  typical  as  those  just  cited,  nevertheless 
a  careful  and  comprehensive  study  of  all  the  elements 
shows  that  in  most  cases  their  properties  vary  periodically 
with  the  atomic  weight.  That  is,  at  certain  regular 
intervals  or  periods  in  the  long  series  elements  will  be 
found  which  have  similar  properties;  in  other  words,  a 
certain  increase  in  atomic  weight  causes  a  reappearance 
or  return  of  properties.  MendelejefT  summarized  these 
facts  in  the  Periodic  Law,  thus :  — 

The  properties  of  the  elements  are  periodic  functions  of 
their  atomic  weights. 

The  term  function  as  used  here  means  the  exhibition  of 
some  special  relation,  viz.  that  of  properties  to  atomic 
weight.  It  is  believed  now  that  the  relation  empha- 
sized by  MendelejefT  is  not  sufficiently  accurate  to  be 
called  a  law.  However,  there  is  undoubtedly  some  rela- 
tion, and  we  are  justified  in  concluding:  (i)  properties 
and  atomic  weight  are  related;  and  (2)  this  relation  is 
exhibited  in  very  many  instances  at  regular  intervals. 

316.  Defects  of  the  Periodic  Table.  —  Examination  of  the  periodic 
table  shows  imperfections.  For  example,  there  are  gaps.  These 
probably  correspond  to  elements  not  yet  discovered.  Three  such 
gaps  which  were  in  the  original  table  have  been  filled.  When  Men- 
delejefT proposed  his  arrangement,  he  predicted  the  discovery  of  three 
elements  having  definite  properties.  These  elements  —  gallium, 
scandium,  and  germanium  —  have  since  been  discovered  and  now 
occupy  their  predicted  places  in  the  table.  Possibly  other  gaps  will 
be  filled  by  newly  discovered  elements  as  was  done  in  the  case  of  the 
zero  group.  According  to  their  atomic  weights  argon  and  potas- 
sium should  exchange  places;  the  same  is  true  of  iodine  and  tellurium. 
Hydrogen  also  lacks  an  acceptable  place. 


CLASSIFICATION  OF  THE   ELEMENTS         265 

EXERCISES 

1.  State  the  general  physical  properties  of  metals  and  non-metals. 

2.  Memorize  the  names  of  the  important  metals  and  non-metals. 

3.  What    metals    are  related   to   (a)  sodium,    (b)  lead,    (c)  copper? 

4.  As   in  Exercise   3,  what  non-metals  to  (a)  carbon,  (b)  chlorine,  (c) 
nitrogen,  (d)  sulphur? 

6.  Classify  the  following  into  metals  and  non-metals:  aluminium,  Zn, 
Na,  silicon,  Ca,  Cu,  bismuth,  Pb,  Ag,  C,  manganese,  O,  H,  iron,  K,  Au, 
N,  hydrogen,  Cl,  Br,  boron,  I,  S,  P. 

6.  What  is  meant  by  (a)  period,  (b)  group,  (c)  family? 

7.  Illustrate  the  periodic  classification  by  two  periods. 

8.  Commit  to  memory  the  names   of  the  elements   in   the   following 
families:    (a)  Argon,  (b)  sodium,  (c)  copper,  (d)  calcium,  (e)  zinc,  (/)  nitro- 
gen, (g)  chlorine,  (h)  iron. 

PROBLEMS  (Review) 

1.  If  36  gm.  of  copper  are  heated  in  air  until  there  is  no  farther  increase 
in  weight,  how  many  grams  will  be  gained? 

2.  Equal  weights  of  sodium  and  calcium  interact  with  water,  and  the 
liberated  gas  is  collected.     Which  metal  yields  the  larger  volume? 

3.  How  many  grams  of  zinc  must  be  used  with  hydrochloric  acid  to 
produce  750  cc.  of  dry  hydrogen  at  20°  C.  and  765  mm.? 

4.  At  what  temperature  will  i  1.  of  chlorine  weigh  the  same  as  i  1.  of 
hydrogen?     (Assume  constant  pressure.) 

6.  Calculate  the  per  cent  of  oxygen  in  (a)  water,  (b)  potassium  chlorate, 
(c)  nitric  acid,  (d)  lime,  (e)  silica,  (/)  borax,  (g)  sulphurous  acid. 

6.  A  candle  in  burning  forms  13.21  gm.  of  carbon  dioxide  and  5.58  gm. 
of  water.     How  much  weight  did  the  candle  lose?     What  volume  of  oxy- 
gen at  o°  C.  and  760  mm.  was  required? 

7.  What  volume  of  air  (free  from  carbon  dioxide  and  water  vapor)  con- 
tains i  gm.  of  nitrogen? 

8.  What  weight  of  sulphur  is  contained  in  500  cc.  of  SO2? 

9.  Suppose  50  1.  of  nitrous  oxide  are  decomposed  into  nitrogen  and 
oxygen.     How  many  volumes  of  the  products  are  formed? 

10.  One  gram  of  a  metal  liberated  1.242  1.  of  hydrogen  from  hydro- 
chloric acid.     The  specific  heat  of  the  metal  was  found  to  be  about  .23. 
Calculate  (a)  the  equivalent  weight  and  (b)  atomic  weight  of  the  metal. 
What  is  the  valence  of  the  metal?     (Standard  conditions.) 

11.  A  compound  has  the  composition  C  =  39.9,  H  =  6.7,  O  =  53.4, 
and  the  vapor  density  is  1.906.     What  is  the  molecular  formula? 


CHAPTER  XXII 

FLUORINE  —  BROMINE  —  IODINE 

Fluorine,  bromine,  and  iodine,  together  with  chlorine, 
constitute  a  family  of  related  elements  often  called  the 
halogens.  The  elements  and  their  analogous  compounds 
have  similar  properties,  differing  mainly  in  degree  (331). 


Fluorine 

317.  Occurrence.  —  Fluorine  is  the  most  active  of 
these  elements,  and  is  never  found  free  in  nature.  It 
occurs  abundantly  in  combination  with  calcium  as  cal- 
cium fluoride  (fluor  spar,  fluorite,  CaF2).  Other  native 
compounds  are  cryolite  (NasAlFe)  and  apatite  (Ca5F 
(PO^a).  Traces  of  fluorine  com- 
pounds are  found  in  bones  and  blood, 
in  the  enamel  of  the  teeth,  and  in 
sea  and  some  mineral  waters. 


318.  Preparation.  —  Fluorine  was  first 
isolated  in  1886  by  the  French  chemist 
Moissan  by  the  electrolysis  of  hydrofluoric 
acid.  The  experiment  was  difficult  and 
dangerous  owing  to  the  corrosive  properties 
of  both  acid  and  element.  Fluorine  is  sel- 
dom prepared  in  the  laboratory. 

A  sketch  of  the  apparatus  used  by  Moissan 
is  shown  in  Fig.  74.  The  U-tube  of  platinum 
has  two  stoppers  of  fluor  spar  (S,  S). 


Fig.  74. — Apparatus  for 
Preparing  Fluorine. 


Through  the  stoppers  pass  the  electrodes  (E,  E)  of  platinum-iridium 
held  in  place  by  screw  caps  (C,  C).  Side  tubes  (T,  T)  allow  the  lib- 
erated gases  (fluorine  and  hydrogen)  to  be  drawn  off.  Hydrofluoric 


FLUORINE  —  BROMINE  —  IODINE  267 

acid  free  from  water  was  put  into  the  U-tube,  and  dry  acid  potassium 
fluoride  (HKF2)  was  added  to  make  the  solution  a  conductor  of 
electricity  —  dry,  liquid  hydrofluoric  acid  itself  being  a  non-conductor. 
The  U-tube  was  then  cooled  to  a  low  temperature  (-23  to  -50°  C.), 
and  on  passing  a  current  through  the  solution  fluorine  was  evolved 
at  the  positive  electrode  (anode)  and  hydrogen  at  the  negative  elec- 
trode (cathode).  The  fluorine  freed  from  the  hydrofluoric  acid  vapor 
was  collected  by  Moissan  at  first  in  a  platinum  tube  with  a  thin  fluor 
spar  plate  closing  each  end,  so  he  could  look  inside  and  examine  the 
gas.  Later  he  found  the  electrolysis  could  be  performed  in  a  copper 
U-tube  and  pure  dry  fluorine  could  be  collected  in  a  glass  tube. 

319.  Properties.  —  Fluorine  is  a  gas  having  a  greenish  yellow 
color,  though  lighter  and  more  yellowish  than  chlorine.     Chemically 
fluorine  is  intensely  active.     Most  elements  unite  with  it  readily, 
the  combining  being  accompanied  by  much  heat  and  light.     The 
compounds  formed  are  fluorides.    It  does  not  combine  with  oxygen 
or  nitrogen,  while  some  metals,  e.g.  gold,  platinum,  and  copper,  are 
not  readily  (or  only  slightly)  attacked  by  it.'    Fluorine  like  chlorine 
withdraws  hydrogen  from  compounds.     (Compare  63,  85.) 

320.  Hydrogen  Fluoride  is  prepared  by  the  interaction 
of  a  fluoride  and  concentrated  sulphuric  acid.     Calcium 
fluoride  is  usually  used,  and  the  experiment  is  performed 
in  a  lead  dish.     The  equation  for  the  reaction  is:  - 

CaF2      +      H2SO4     =       2HF     +    CaS04 

Calcium  Sulphuric  Hydrogen  Calcium 

Fluoride  Acid  Fluoride  Sulphate 

Hydrogen  fluoride  is  a  colorless  liquid  which  boils  at 
about  19°  C.  It  is  very  volatile,  and  the  gas  forms  fumes 
in  air  and  dissolves  readily  in  water.  The  solution  is  the 
commercial  hydrofluoric  acid.  Hydrogen  fluoride  in  the 
form  of  gas,  liquid,  or  solution  is  a  dangerous  substance. 
The  gas  is  extremely  poisonous,  and  the  liquid,  if  dropped 
on  the  skin,  produces  terrible  sores.  Owing  to  its  corro- 
sive properties,  hydrofluoric  acid  is  kept  in  hard  rubber  or 
wax  bottles.  Hydrofluoric  acid  has  chemical  properties 


268 


CHEMISTRY 


similar   to   hydrochloric   acid.     Thus,   it   interacts   with 
metals  forming  hydrogen  and  fluorides  and  it  neutralizes 
bases    forming    salts   and   water.     Unlike 
hydrochloric   acid,  it   forms  both  normal 
and  acid  fluorides,  e.g.  potassium  fluoride 
(KF)  and  acid  potassium  fluoride  (HKF2). 
The    acid    and    the   moist   gas    attack 
glass  and  are  used  extensively  in  etching. 
Glass  as  we  have  already  seen  is  a  mix- 
ture of  silicates  (309).     Hydrofluoric  acid 
interacts   with    these    silicates   and   forms 
among   other   substances   a  volatile  com- 
ing  on    Glass    P°und  called  silicon    tetrafluoride    (SiF4). 
Tumbler — De-    Thus    the    acid    disintegrates    the    glass 
signed  and  Ex-    —  literally    "eats"    or    etches   it.     Typi- 

ecuted     by     a         ,  , .  r 

Pu  jj  cal    equations    for    the   reactions    in    the 

case  of  ordinary  sodium  glass  are :  - 


Na2SiO3 

Sodium 
Silicate 

CaSiO3 

Calcium 
Silicate 


6HF     =      SiF4     +  2NaF     +  3H2O 

Hydrofluoric          Sodium          Sodium  Water 
Acid           Tetrafluoride     Fluoride 

6HF     =     SiF4     +       CaF2     +  3H2O 

Hydrofluoric         Silicon                 Calcium  Water 
Acid          Tetrafluoride           Fluoride 


Hydrofluoric  acid  also  interacts  with  silicon  dioxide,  the 
equation  for  the  reaction  being  :  - 


Si02  +      4HF       =      SiF4     + 

Silicon         Hydrofluoric  Silicon 

Dioxide  Acid  Tetrafluoride 


Water 


In  etching  with  hydrofluoric  acid,  the  glass  is  thinly  coated  with 
wax,  and  the  design  or  scale  marks  to  be  etched  are  scratched  through 
the  wax.  The  glass  is  then  exposed  to  the  gas  or  liquid,  which  attacks 
the  unprotected  places.  When  the  wax  is  removed,  a  permanent 


FLUORINE  —  BROMINE  —  IODINE 


269 


etching  is  left.  Sometimes  the  design  or  marking  is  made  more  con- 
spicuous by  filling  the  etched  cavity  with  an  insoluble  white  or  black 
substance.  Hydrofluoric  acid  is  utilized  in  marking  the  scales  on  the 
thermometers,  tubes,  and  other  graduated  glass  instruments,  and  also 
in  etching  designs  on  glassware  (Fig.  75).  (See  Part  II,  Exp.  152.) 


Bromine 

321.  Occurrence.  —  Uncombined    bromine    is    never 
found,  but  bromides  are  widely  distributed,  especially  mag- 
nesium, sodium,  potassium,  and  calcium  bromides.     The 
salt  springs  and  wells  of  Ohio,  West  Vir- 

ginia, Pennsylvania,  and  Michigan  con- 
tain bromides,  and  large  quantities  are 
found  in  the  salt  deposits  at  Stassfurt 
(377). 

322.  Preparation.  --  Bromine    is   pre- 
pared    in     the    laboratory    by    heating 
potassium  bromide  with    manganese    di- 
oxide and  sulphuric   acid.     If   an   appa- 
ratus like  that  shown  in  Fig.  76  is  used, 

part  of  the  bromine  vapor   escapes    and    Flg-  ?6'      APPa~ 

.       .  .  ratus  for  Prepar- 

part    condenses   to   a   liquid   which   col-       ing  Bromine. 
lects    in    the    V-bend.      The     complete 
equation  for  the  reaction,  which  takes  place  in  several 
stages,  is:  - 


2KBr 

Potassium        Sulphuric 
Bromide  Acid 


Mn02  =    Br2   +  MnSO4 

Manganese      Bromine      Manganese 
Dioxide  Sulphate 


2HKS04  +  2H2O 

Acid  Potas-         Water 
sium  Sulphate 


Another  method  consists  in  warming  a  bromide  solution 
with  chlorine  water;   an  equation  for  this  method  is:  - 


MgBr2 

Magnesium  Bromide 


C12     =     Br2 

Chlorine        Bromine 


MgCl2 

Magnesium  Chloride 


270  CHEMISTRY 

In  one  commercial  process  the  salt  water  containing  bromides  is 
subjected  to  electrolysis;  bromine  and  some  chlorine  are  liberated, 
but  the  chlorine  assists  the  process  by  displacing  bromine  from  the 
bromides  (see  last  equation  in  the  preceding  paragraph).  In  another 
process  sulphuric  acid  and  potassium  chlorate  are  heated  with  a  con- 
centrated solution  of  bromides  called  bittern,  which  is  left  after  the 
sodium  chloride  has  been  removed  by  crystallization  from  the  salt 
water;  sometimes  the  bittern  is  treated  with  chlorine.  It  is  inter- 
esting to  note  that  the  French  chemist  Balard,  who  discovered  bro- 
mine in  1826,  obtained  it  from  bittern.  (See  Part  II,  Exp.  164.) 

323.  Properties.  —  Bromine  is  a  reddish  brown  liquid 
which  is  about  three  times  as  heavy  as  water.     It  is  a 
volatile  liquid,  boiling  at  about  59°  C.     The  vapor,  which 
is  given  off  freely,  has  a  disagreeable  odor.     This  prop- 
erty suggested  the  name  bromine   (from  a  Greek  word 
meaning    a    stench).     The    vapor    irritates    the    mucous 
membrane  of  the  eyes,  nose,   and   throat;    a  bottle  of 
bromine  should  not  be  opened  unless  it  is  in  the  hood. 
Liquid  bromine  burns  the  flesh  frightfully,  and  the  utmost 
care  should  be  used  in  preparing  or  working  with  this  sub- 
stance.    Bromine  is  somewhat  soluble  in  water.     The 
solution,  called  bromine  water,  has  a  reddish  brown  color. 
Bromine   also   dissolves   in   carbon   disulphide,   and   the 
solution  is  reddish  yellow.      (See  Part  II,  Exp.  153.) 

Many  chemical  properties  of  bromine  are  similar  to 
those  of  chlorine,  though  bromine  is  less  active.  Thus,  it 
combines  with  metals  and  other  elements;  it  also  bleaches. 

324.  Uses.  —  Bromine  is  used  to  prepare  bromides  and  certain 
dyes  tuffs. 

325.  Compounds  of  Bromine  are   similar   to   those   of   chlorine. 
Hydrogen  bromide  (HBr)  is  a  colorless,  pungent  gas,  which  fumes 
in  the  air  and  dissolves  freely  in  water,  forming  a  solution  called 
hydrobromic    acid.      Its   other    properties    closely   resemble    .those 
of   hydrochloric   acid.      Bromides   are   salts   of  hydrobromic   acid, 
though    many  are  formed    by  direct  combination  with  bromine. 


FLUORINE  —  BROMINE  —  IODINE  271 

Potassium  bromide  (KBr)  is  a  white  solid,  made  by  decomposing 
iron  bromide  with  potassium  carbonate;  it  is  used  as  a  medicine. 
Silver  bromide  (AgBr)  is  used  extensively  in  photography. 

Iodine 

326.  Occurrence.  —  Iodine,  like  chlorine  and  bromine, 
is  found  in  nature  only  in  compounds.     They  are  widely 
distributed,  though  the  quantity  in  a  single  place  is  usually 
small.     Tobacco,  water  cress,  cod-liver  oil,  oysters,  and 
sponges  contain  minute  quantities.     Sea  water  contains 
a   very   small  proportion   of   combined  iodine.     This  is 
assimilated  by   certain  seaweeds,   and  can    be  obtained 
from  their  ashes.     Much  iodine  was  formerly  extracted 
from  seaweed.     The  French  chemist  Courtois,  who  first 
isolated  iodine  in  1812,  obtained  it  from  seaweed.     Iodine 
compounds,  chiefly  sodium  iodate  (NaI03),  occur  in  the 
deposits  of  saltpeter  (sodium  nitrate)  in  Chile,  and  most 
of    the    iodine  of    commerce  is  now  obtained  from  this 
source. 

327.  Preparation.  —  Iodine   is   prepared  in  the  labo- 
ratory by  a  method  similar  to  that  used  for  bromine. 
Potassium  iodide,  manganese  dioxide,  and  sulphuric  acid 
are  heated  together.     The  iodine  is  evolved  as  a  violet 
colored  vapor,  which  condenses  on  the  colder  part  of  the 
vessel   in    dark   purplish   gray    crystals.     The   complete 
equation  for  the  chemical  changes  involved  is:  - 

2KI    +  3H2SO4  +  Mn02     =    I2    +  MnSO4  +  2HKSO4  +  2H2O 

Potassium         Sulphuric        Manganese        Iodine        Manganese        Acid  Potas-        Water 
Iodide  Acid  Dioxide  Sulphate         sium  Sulphate 

Iodine  is  prepared  on  a  commercial  scale  from  crude  Chile  saltpeter 
or  from  seaweed.  In  the  first  process  sodium  sulphites  are  added  to 
the  liquid  left  after  the  sodium  nitrate  has  been  removed  by  crystal- 
lization from  a  solution  of  the  crude  saltpeter.  Iodine  is  precipitated 
according  to  the  equation:  — 


272  CHEMISTRY 

2NaI03  +  3Na2SO3  +  2HNaSO3   =    I,    +   sNa£O4   +  H2O 

Sodium  Sodium  Acid  Sodium         Iodine  Sodium  Water 

lodate  Sulphite  Sulphite  Sulphate 

In  the  second  process  the  seaweed,  found  principally  on  the  shores 
of  France,  Scotland,  and  Norway,  is  allowed  to  ferment  and  dry,  and 
is  then  burned.  From  the  ash  the  soluble  salts  are  extracted,  and 
from  the  purified  solution,  in  which  the  iodides  are  dissolved,  the 
iodine  is  obtained  by  treatment  with  sulphuric  acid  or  with  chlorine. 

328.  Properties.  —  Iodine  is  a  dark  grayish  crystal- 
line solid,  which  is  nearly  five  times  as  heavy  as  water. 
It  is  volatile  at  ordinary  temperatures,  and  when  gently 
heated  changes  into  a  beautiful  violet  colored  vapor, 
which  readily  solidifies  on  a  cold  surface,  often  in  the  upper 
part  of  the  test  tube  in  which  the  iodine  is  heated.  This 
property  of  iodine,  viz.  ready  transformation  from  solid 
into  vapor  and  back  to  solid,  is  utilized  in  purifying 
iodine.  The  crude  substance  is  heated  gently  and 
the  vapor  is  condensed  in  a  series  of  vessels;  the  non- 
volatile impurities  remain  behind.  This  process  is  called 
sublimation  and  is  frequently  used  to  purify  substances. 
The  striking  color  of  the  vapor  suggested  the  name  iodine 
(from  a  Greek  word  meaning  violetlike),  which  was  given 
to  the  element  by  the  English  chemist  Davy,  who  studied 
the  substance  soon  after  its  discovery.  The  vapor  is 
nearly  nine  times  as  heavy  as  air,  and  has  an  odor  resem- 
bling dilute  chlorine,  though  less  irritating.  Iodine  melts 
at  about  114°  C.  and  boils  at  about  185°  C.  Iodine  turns 
cold  starch  solution  blue.  The  presence  of  a  minute 
trace  of  iodine  may  be  thus  detected.  The  exact  nature 
of  this  blue  substance  is  unknown.  The  presence  of 
starch  in  many  vegetable  substances  can  be  shown  by 
this  delicate  test.  Iodine  dissolves  slightly  in  water  and 
freely  in  alcohol,  chloroform,  carbon  disulphide,  ether, 
and  potassium  iodide  solution.  The  chloroform  and  car- 


FLUORINE  —  BROMINE  —  IODINE  273 

bon  disulphide  solutions  are  violet,  but  the  others  are 
brown,  or  even  black.  Iodine  and  its  solutions  turn  the 
skin  brown.  (See  Part  II,  Exp.  153.) 

The  chemical  properties  of  iodine  resemble  those  of  chlo- 
rine and  bromine,  but  iodine  is  less  active.  Bromine  and 
chlorine  displace  iodine  from  many  of  its  compounds. 

329.  Uses.  —  A  solution  of  iodine  in  alcohol  (or  in  alcohol  and 
potassium  iodide),  called  tincture  of  iodine,  is  used  as  an  application 
for  the  skin  to  prevent  the  spread  of  eruptions  or  to  reduce  swellings. 
Iodine  is  used  to  make  medicinal  preparations,  especially  iodoform 
(CHI3),  which  is  used  as  an  antiseptic  for  wounds.     Large  quantities 
of  iodine  are  used  in  making  iodides  and  certain  drugs  and  dyes. 

330.  Compounds  of  Iodine  resemble  the  corresponding  ones  of 
chlorine  and  bromine,  though  some  are  less  stable.     Hydriodic   acid 
(HI)  is  much  like  hydrobromic  and  hydrochloric  acids,  though  unlike 
them  in  being  a  reducing  agent.     Iodides  are  salts  of  hydriodic  acid. 
In  general  behavior  they  are  similar  to  bromides  and  chlorides.      The 
best  known  salt  is  potassium  iodide  (KI). 

331.  The  Halogen  Elements  and  the  Periodic  Classi- 
fication. —  The  halogen  elements  furnish  a  typical  illus- 
tration of  the  periodic  classification.  These  elements,  as 
arranged  in  the  periodic  table,  increase  in  atomic  weight 
from  fluorine  (19.0)  through  chlorine  (35*46)  and  bromine 
(79.92)  to  iodine  (126.92),  and  many  of  their  properties 
are  graded  in  this  order.  Thus,  as  we  pass  from  fluorine 
to  iodine  the  specific  gravity  increases,  the  color  grows 
deeper,  the  volatility  decreases,  and  the  melting  points 
of  the  solidified  elements  and  the  boiling  points  of  the 
liquefied  elements  increase.  The  intensity  of  the  chemi- 
cal action  decreases  as  we  pass  from  fluorine  to  iodine. 
Other  properties  of  the  elements  and  many  properties  of 
their  compounds  emphasize  the  fundamental  principle  of 
the  periodic  classification,  viz.  properties  are  a  periodic 
function  of  atomic  weights. 


274  CHEMISTRY 

EXERCISES 

1.  Review  topics:    (a)  Chlorine,  (6)  compounds  of  fluorine  and  silicon, 
(c)  periodic  classification  of  the  elements. 

2.  Summarize  the  chief  properties  of  fluorine  and  hydrogen  fluoride. 

3.  Describe  the  process  of  etching  glass.     Express  the  essential  change 
by  an  equation. 

4.  Give  the  equations  for  the  preparation  of  (a)  hydrogen  fluoride,  (&) 
bromine,  (c)  iodine,  (d)  silver  bromide. 

6.  Practical  questions:  (i)  Why  did  Moissan  use  acid  potassium  fluo- 
ride in  the  electrolysis  of  hydrofluoric  acid?  (2)  Hydriodic  acid  and  iodine 
compounds  often  turn  dark  on  standing.  Why?  (3)  What  gas  resembles 
bromine  vapor  in  color?  (4)  How  does  bromine  differ  from  all  other  ele- 
ments previously  studied  (in  this  book)?  (5)  How  would  you  identify  by 
experiment  (a)  sand,  (b)  tincture  of  iodine,  (c)  calcium  fluoride,  (d)  potas- 
sium bromide,  (e)  silver  iodide?  (6)  How  would  you  distinguish  a  chloride 
from  an  iodide? 

6.  Compare  the  properties  of  fluorine,  chlorine,  bromine,  and  iodine. 

7.  Essay  topics:    (a)  Discovery  of    the    halogen  elements.     (&)  Iodine 
industry  in  Chile,     (c)  Etching  with  hydrofluoric  acid,     (d)  Uses  of  iodine. 

PROBLEMS 

1.  Calculate  the  per  cent  of  fluorine  in  (o)  hydrogen  fluoride  (HF), 
(b)  silicon  tetrafluoride,  (c)  apatite,  (d)  calcium  fluoride. 

2.  Calculate  the  per  cent  of  bromine  or  iodine  in  (a)  sodium  bromide, 
(6)  hydrogen  bromide,    (c)  calcium  iodide,    (d)  sodium  iodate,    (e)  hydro- 
gen iodide,  (/)  iodoform. 

3.  How   much    (a)  calcium   sulphate .  and    (b)  hydrogen    fluoride   are 
formed  by  heating  60  gm.  of  fluor  spar  with  sulphuric  acid? 

4.  How  much  iron  iodide  (Feals)  can  be  made  by  the  interaction  of 
iron  and  300  gm.  of  iodine? 

5.  How  much  potassium  bromide  (75  per  cent  pure)  is  necessary  to 
prepare  47  gm.  of  bromine? 

6.  How  much  potassium  iodide  (80  per  cent  pure)  is  necessary  to  pre- 
pare 25  gm.  of  iodine? 

7.  Write  the  formulas  of  the  fluoride,  bromide,  and  iodide  of  Al,  am- 
monium, Ba,  Ca,   copper   (ous  and  ic),  Fe,n  Fem,  Pb,  magnesium,  Sbnl, 
Si,  Hg  (ous  and  ic),  Sn11,  SnIV,  zinc. 

8.  Calculate  the  atomic  weight  of  fluorine,  bromine,  or  iodine  from 
the  following:    (a)  i  gm.  of  CaF2  gives  1.745  gm.  of  CaSO4;    (b)  3.946  gm. 
of  Ag  (dissolved   in  HNO3)  require  4.353  gm.  of  KBr  for  precipitation; 
(c)  6.3835  gm.  of  silver  iodide  give  3.8965  gm.  of  silver  chloride. 


CHAPTER   XXIII 
PHOSPHORUS  —  ARSENIC  —  ANTIMONY  —  BISMUTH 

Phosphorus 

332.  Occurrence.  —  Free  phosphorus  is  not  found  in 
nature,  but  phosphorus  compounds  are  numerous  and 
some,    especially    those    related    to    calcium    phosphate, 
are  abundant.     The  most   common   phosphate  is  apatite 
(Ca5F(P04)3).     Small  amounts  of  phosphates  are  present 
in  all  fertile  soils  and  in  many  iron  ores;  bones  are  about 
80  per  cent  calcium    phosphate.     Complex   phosphorus 
compounds  are  essential  components   of    the    germs  of 
seeds  and  of  the  nerves,  brain,  blood,  and  muscles  of 
animals.     (See  344.) 

333.  Preparation.  —  Phosphorus  is  manufactured  from  bone  ash 
or  from  native  phosphates.     In  the  older  process  the  finely  ground 
material  is  mixed  with  enough  sulphuric  acid  to  produce  the  follow- 
ing change:  — 


Ca3(P04)2        +    3H2SO4     =     2H3P04     + 

Calcium  Sulphuric  Phosphoric  Acid  Calcium 

Phosphate  Acid  (Ortho-)  Sulphate 

The  phosphoric  acid  is  mixed  with  sawdust,  coke,  or  charcoal,  and 
dried,  being  changed  thereby  according  to  the  equation  — 

H3P04  =  HP03  +        H20 

Phosphoric  Acid  (Ortho-)        Phosphoric  Acid  (Meta-) 

The  dried  mass  is  heated  to  a  high  temperature  in  clay  retorts,  the 
change  thus  produced  being  substantially  — 

4HPO3     •+       i2C       =       P4       +       2H2       +       I2CO 

Phosphoric  Acid  (Meta-)     Carbon  Phosphorus  Hydrogen        Carbon  Monoxide 


276 


CHEMISTRY 


The  phosphorus  distils  as  a  vapor  through  a  pipe  into  a  trough  of 
water  where  it  condenses  as  a  heavy  liquid.  In  the  new  process 
phosphorus  is  manufactured  in  an  electric  furnace.  The  mixture  of 
phosphate,  carbon,  and  sand  is  introduced  at  A  and  fed  into  the 
furnace  by  the  screw  B.  An  electric  current  passed  between  the  elec- 
trodes E,  E  produces  the  intense  heat  needed  for  the  chemical  change. 
The  phosphorus  vapor  escapes  through  C  into  a  condenser;  the  liquid 
residue,  which  is  essentially  calcium  silicate,  is  drawn  off  as  slag  at 
D  (Fig.  77).  The  equation  for  the  chemical  change  is  — 

2Ca3(PO4)2   +  6SiO2    +    loC    =  P4   +    170)   +     6CaSi03 


Calcium 
Phosphate 


Sand 


Carbon    Phosphorus      Carbon 
Monoxide 


Calcium 
Silicate 


C 


The  product  obtained  by  either  method  is  purified  and  pressed  through 
chamois  skin  or  canvas  into  molds  cooled  by  water. 

334.    Properties.  —  Phosphorus    exists    in    two    allo- 

tropic  modifications,  called  re- 
spectively yellow,  white,  or 
waxy  phosphorus  and  red  or 
amorphous  phosphorus  (see 
181).  The  phosphorus  pre- 
pared by  the  methods  just  de- 
scribed is  the  yellow  or  ordinary 
form.  It  is  a  colorless  or 
slightly  yellow,  translucent 
solid.  The  color  deepens  by 
exposure  to  light.  At  ordinary 
temperatures  it  is  like  wax, 
but  at  low  temperatures  it  is 
brittle.  It  melts  at  44°  C.  but 
it  should  be  melted  under  water.  When  exposed  to 
air  it  soon  gives  off  white  fumes,  and  at  about  34°  C. 
takes  fire  and  burns  with  a  brilliant  flame.  Hence 
phosphorus  is  kept  beneath  water  and  should  not  be 
handled  unless  it  is  wet;  indeed  it  is  better  not  to 
touch  it  at  all,  but  to  use  wet  forceps  when  it  is  necessary 


Fig.  77.  —  Electric  Furnace  for 
the  Manufacture  of  Phospho- 
rus. 


PHOSPHORUS  277 

to  transfer  it  or  to  hold  it  while  it  is  being  cut  under 
water.  Moreover  unusual  care  should  be  taken  not  to 
leave  bits  of  phosphorus  in  deflagrating  spoons  or  lying 
about  the  laboratory.  In  moist  air  it  is  slightly  luminous, 
as  may  be  easily  seen  by  rubbing  the  head  of  a  phos- 
phorus tipped  match  in  a  dark  room.  This  property  gave 
the  element  its  name  (from  a  Greek  word  meaning  light 
bringer).  The  ease  with  which  it  ignites  makes  phos- 
phorus dangerous  to  handle.  Burns  from  it  heal  slowly. 
It  is  very  poisonous,  and  the  fumes  cause  a  dreadful  disease, 
which  rots  the  bones,  especially  the  jaw  bones.  It  is 
practically  insoluble  in  water,  but  dissolves  readily  in  car- 
bon disulphide  and  slightly  in  sodium  hydroxide  solution. 
Yellow  phosphorus  has  a  faint  odor,  which  may  be  easily 
detected  by  smelling  a  match  head.  Red  phosphorus  is 
made  by  heating  ordinary  phosphorus  to  25o°-3OO°  C.  in  a 
closed  vessel  freed  from  air.  Red  phosphorus  is  a  reddish 
brown  powder.  It  is  opaque  and  odorless,  does  not  glow 
in  the  air,  nor  does  it  ignite  until  heated  to  about  260°  C. 
It  is  not  poisonous,  and  does  not  dissolve  in  carbon  disul- 
phide. Its  specific  gravity  varies  from  2.1  to  2.3,  that  of 
the  yellow  form  being  1.83.  It  does  not  combine  with 
oxygen  at  ordinary  temperatures  and  being  less  danger- 
ous than  yellow  phosphorus  can  be  handled  safely.  (See 
Part  II,  Exp.  164.) 

The  vapor  density  of  both  kinds  of  phosphorus  up  to  about  1500°  C. 
is  such  that  its  molecule  must  contain  four  atoms,  and  the  usual 
molecular  formula  of  the  vapor  is  written  P4. 

335.  Phosphorus  Oxides.  —  There  are  two  important  oxides. 
Phosphorus  trioxide  (P2O3)  is  a  white  solid  formed  by  the  slow  oxida- 
tion of  phosphorus  or  by  burning  phosphorus  in  a  limited  supply  of 
air.  It  has  the  odor  of  phosphorus  and  is  poisonous.  Warmed  in 
the  air,  it  changes  into  the  pentoxide.  It  unites  with  water  to  form 
phosphorous  acid,  thus  — 


278  CHEMISTRY 


P203  +       3H20  2H3P03 

Phosphorus  Trioxide  Water  Phosphorous  Acid 

Phosphorus  pentoxide  (P2O5)  is  the  white,  snowlike  solid  formed  by 
burning  phosphorus  in  an  abundant  supply  of  air.  It  is  very  deliques- 
cent, quickly  withdrawing  moisture  from  air  and  combining  vigorously 
with  water  with  a  hissing  noise.  It  is  often  used  in  the  laboratory  to 
dry  gases,  being  much  more  effective  than  calcium  chloride  and  sul- 
phuric acid,  which  are  commonly  employed. 

336.  Phosphoric  Acids  and  Phosphates.  —  The  three 
phosphoric    acids    are    orthophosphoric    (HsPCX),    meta- 
phosphoric  (HPO3),  and  pyrophosphoric  (H4P2O7).     Each 
acid  forms  many  salts  called  phosphates. 

337.  Orthophosphoric  Acid  is  a  white,  crystalline,  deliquescent 
solid,  though  it  usually  is  sold  as  a  thick  solution.     A  commercial 
grade  is  obtained  as  a  by-product  in  the  first  step  of  the  manufacture 
of  phosphorus  from  calcium  phosphate  (333).     The  pure  acid  is  made 
by  oxidizing  phosphorus  with  nitric  acid,  or  by  dissolving  phosphorus 
pentoxide  in  hot  water,  thus  (in  the  latter  case)  :  — 


P2O6  +      3H20  2H3PO4 

Phosphorus  Pentoxide  Water  Orthophosphoric  Acid 

338.  Metaphosphoric  Acid  at  ordinary  temperatures  is  a  glassy 
solid,    and  is   therefore  often  called  glacial  phosphoric  acid.     It  is 
formed  by  heating  orthophosphoric  acid,  thus:  — 

H3P04  =  HP03  +        H20 

Orthophosphoric  Acid  Metaphosphoric  Acid  Water 

It  may  be  formed  by  dissolving  the  pentoxide  in  cold  water,  thus:  — 
P205  +  H20  =  2HP03 

Metaphosphoric  acid  dissolves  readily  in  water,  and  the  solution 
changes  into  orthophosphoric  acid  —  slowly  in  the  cold,  rapidly  when 
boiled. 

339.  Pyrophosphoric  Acid  is  an  amorphous,  glassy  (but  sometimes 
crystalline)  solid.     It  is  formed  by  heating  orthophosphoric  acid  to 
about  215°  C.     The  equation  for  the  reaction  is  — 

2H3P04  H4P207        +        H2O 

Orthophosphoric  Acid  Pyrophosphoric  Acid 


PHOSPHORUS  279 

340.  Phosphates  are  salts  of  the  acids  just  discussed. 
Orthophosphoric  acid  is  tribasic  and  hence  forms  three 
series  of  salts.    Orthophosphates  —  usually  called  simply 
phosphates  —  have  several  names  based  on  the  replace- 
ment of  the  hydrogen.     Thus,   Na3PO4  is  normal,   tri- 
sodium,    tertiary    phosphate;     HNa2PO4    is    disodium, 
secondary  phosphate;   and  H2NaPO4  is  monosodium,  pri- 
mary phosphate.     Disodium  phosphate  (HNa2PO4)  is  the 
commercial    " sodium    phosphate."       The    "acid    phos- 
phate" sold  as  a  beverage  is  a  solution  of  one  or  more 
acid  calcium  phosphates  (HCaPO4  and  H4Ca(PO4)2).     In 
phosphate  baking  powder   the   acid  ingredient  is  mono- 
calcium   phosphate    (compare    243,   269).      With   silver 
nitrate,    orthophosphoric    acid    and    soluble    orthophos- 
phates  precipitate  yellowish  silver  phosphate  (Ag3P04); 
they  also  precipitate  yellow  ammonium  phosphomolyb- 
date  from  an  excess  of  a  nitric  acid  solution  of  ammon- 
ium   molybdate.     These    reactions    serve    as    tests    for 
orthophosphoric  acid  and  its  salts.      (See  Part  II,  Exps. 
159,  160,  113  E.)     Metaphosphates  are  formed  by  heat- 
ing primary  (or  mono-)  phosphates,  thus :  - 

H2NaP04  NaP03  +     H2O 

Monosodium  Phosphate        Sodium  Metaphosphate 

Pyrophosphates  are  formed  by  heating  secondary  (or  di-) 
phosphates,  thus:- 

2HNa2P04      =          Na4P207         +     H2O 

Disodium  Phosphate        Sodium  Pyrophosphate 

341.  Hypophosphites  are  produced  by  treating  phosphorus  with 
an  alkali.     They  are  often  used  as  medicines. 

342.  Other   Compounds   of  Phosphorus.  —  Phosphine   (PH3)   is 
analogous  to  ammonia  (NH3),  though  it  is  not  alkaline.     Phosphorus 
trichloride  (PC13)  is  a  disagreeable  smelling  liquid,  made  by  the  com- 
bustion of  dry  chlorine  and  phosphorus;  and  phosphorus  pentachloride 


280  CHEMISTRY 

(PC15)  is  a  greenish  solid  made  by  passing  chlorine  into  a  vessel  con- 
taining the  trichloride. 

343.  Matches.  —  Phosphorus  until  recently  was  chiefly  used  in 
the  manufacture  of  matches.     Now  however  a  phosphorus  sulphide 
(P4S3)  is  generally  used  in  the  United  States  as  a  substitute  for  the 
element.     This  change  was  made  on  account  of  a  prohibitive  tax 
upon  phosphorus  matches.     The  law  imposing  the  tax  was  passed 
partly  on  account  of  the  fires  accidentally  caused  by  ignition  of 
matches  but  mainly  to  protect  the  workmen  from  the  disease  caused 
by  breathing  phosphorus  fumes.     Ordinary  matches  are  made  by 
dipping  one  end  of  the  match  sticks  first  into  melted  sulphur  or 
paraffin  and  then  into  the  "phosphorus  mixture."     The  latter  consists 
usually  of  different  proportions  of  (i)  phosphorus  sulphide,  (2)  man- 
ganese dioxide  or  another  oxidizing  substance,  and  (3)  glue  or  some 
other  binding  material  mixed  with  a  little  coloring  matter.     These 
matches  are  the  ordinary  friction  or  sulphur  kind.     By  rubbing  them 
on  a  rough  surface  the  friction  generates  enough  heat  to  ignite  the 
phosphorus,  which  continues  to  burn  owing  to  the  oxygen  supplied 
(mainly)  by  the  oxidizing  agent,  and  the  heat  thereby  produced  sets 
fire  to  the  sulphur  or  paraffin,  and  this  in  turn  kindles  the  wood. 
In  safety  matches  the  head  is  usually  a  colored  mixture  of  antimony 
sulphide,  potassium  chlorate,  and  glue,  while  the  surface  on  the  box 
upon  which  the  match  must  be  rubbed  to  ignite  is  a  mixture  of  red 
phosphorus,  glue,  and  powdered  glass. 

344.  Relation  of  Phosphorus  to  Life.  —  Phosphorus  is 
essential  to  the  growth  of  plants  and  animals.     Plants 
take  phosphates  from  the  soil  and  store  up  the  phosphorus 
compounds,  especially  in  their  seeds.     Animals  eat  this 
vegetable  matter,  assimilate  the  phosphorus  compounds, 
and  deposit  them  in  the  bones,  brain,  and  nerve  tissue. 
Most  of  these  phosphorus  compounds  are  complex.     Bones 
however,  as  already  stated,  consist  of  about  80  per  cent 
of  calcium  phosphate   (Ca3(PO4)2).     The  complex  phos- 
phorus compounds  consumed  by  animals  as  parts  of  food 
are  transformed  by  vital  processes  into  phosphates  which 
are  eliminated  and  thus  often  find  their  way  back  into  the 


PHOSPHORUS  281 

soil  to  some  extent.  Here  they  are  taken  up  again  by 
plants,  converted  into  complex  compounds,  stored  up  in 
the  tissues  and  seeds,  which  are  in  turn  eaten  by  animals. 
And  so  the  process  goes  on  —  a  phosphorus  cycle  analo- 
gous to  the  carbon  cycle  (186). 

In  order  to  furnish  plants  with  phosphorus  various  phosphorus- 
bearing  substances  are  added  to  the  soil  in  the  form  of  natural  or 
artificial  fertilizers.  Natural  fertilizers  are  (i)  stable  refuse,  which 
always  contains  some  of  the  .phosphates  from  the  food  originally 
fed  to  the  animals;  (2)  guano,  which  is  dried  excrement  and  carcasses 
of  the  sea  birds  that  once  lived  in  vast  numbers  in  Peru  and  Chile; 
and  (3)  phosphate  slag,  which  is  a  phosphorus  by-product  obtained  in 
manufacturing  steel.  These  natural  phosphatic  materials  as  well  as 
bones  are  ground  and  spread  upon  the  soil.  Artificial  fertilizers  are 
made  from  phosphate  rock.  This  occurs  in  large  beds  in  South  Caro- 
lina, Tennessee,  and  Florida,  which  yield  about  a  million  tons  a  year. 
It  consists  of  the  hardened  remains  of  land  and  marine  animals,  and 
is  mainly  tricalcium  phosphate  (Ca3(PO4)2).  It  is  insoluble  in  water, 
and  must  be  changed  into  the  soluble  monocalcium  salt  (H4Ca(PO4)2), 
so  that  it  can  be  evenly  distributed  through  the  soil  and  easily  taken 
up  by  plants.  This  soluble  salt  is  called  "superphosphate  of  lime." 
When  phosphate  rock  is  treated  with  sulphuric  acid,  the  changes 
involved  may  be  written  thus:  — 

Ca3(P04)2     +     2H2SO4     =     H4Ca(P04)2     +     2CaSO4 

Tricalcium  "  Superphosphate  Calcium 

Phosphate  of  Lime "  Sulphate 

Ca3(PO4)2     +     3H2SO4     =       2H3PO4         +    3CaSO4 

Phosphoric  Acid 

Ca3(P04)2     +    H2SO4      =     H2Ca2(P04)2    +     CaSO4 

Dicalcium  Phosphate 

Sometimes  -"  superphosphate "  is  mixed  with  compounds  of  nitrogen 
and  of  potassium  to  form  a  complete  fertilizer  (98,  386).  The  law 
requires  the  dealer  to  state  the  analysis  of  the  fertilizer  on  the  bag 
or  label.  The  per  cent  of  phosphorus  is  usually  stated  as  per  cent  of 
P205,  which  is  popularly  but  incorrectly  called  "phosphoric  acid." 


282  CHEMISTRY 

Arsenic 

345.  Occurrence.  —  Arsenic  is  occasionally  found  free 
in   nature,  but  it  usually  occurs  combined  with  sulphur 
or   a  metal,   or  with  both,  e.g.  as  realgar   (As2S2),  orpi- 
ment  (As2Sa),  arsenopyrite  or  mispickel  (FeSAs).     Small 
quantities  occur  in  ores,  especially  sulphides. 

346.  Properties  and  Use.  —  Arsenic  is  a  brittle,  steel-gray  solid 
having  a  metallic  luster.     Heated  in  the  air,  it  volatilizes  without 
melting,  and  the  vapor  has  an  odor  like  garlic.     At  about  180°  C. 
it  burns  in  the  air  with  a  bluish  flame,  forming  white  arsenious  oxide 
(As2O3).     Arsenic  is  used  to  harden  the  lead  which  is  made  into  shot. 

The   molecules  of  arsenic  vapor  at  about  650°  C.   contain  four 
atoms,  hence  the  molecular  formula  As4. 

347.  Arsenious  Oxide,  As2O3,  is  the  most  important 
compound  of  arsenic.    It  is  often  called  "  white  arsenic," 
or  simply  "arsenic."     It  is  obtained  by  roasting  arsenic 
ores.    It  is  odorless,  has  a  slight' taste,  and  dissolves  slightly 
in  cold  water.     It  is  converted  readily  by  hot  hydro- 
chloric acid  into  soluble  arsenic  trichloride  (AsCls),  which 
is  a  convenient  solution  of  arsenic  to  use  in  the  labora- 
tory.    Arsenic  trioxide  is  a  rank  poison.     The  antidote  for 
arsenic  poisoning  is  fresh  ferric  hydroxide,  which  is  made 
by  adding  ammonium  hydroxide  to  a  ferric  salt,  e.g.  ferric 
chloride,  which  forms  an  insoluble  substance  with  the 
arsenic  compound.     Arsenic  trioxide  is  used  to  a  limited 
extent  in  making  pigments  for  green  paints,  as  the  poison- 
ous ingredient  of  fly  and  rat  poison,  in  the  manufacture 
of  glass  (especially  plate  and  window  glass),  in  making 
arsenic  compounds  (e.g.  insecticides  (348)),  for  destroy- 
ing weeds,  and  in  preserving  skins  in  museums.     As  a 
medicine  it  is  sometimes  used  to  purify  the  blood. 

348.  Other  Arsenic  Compounds.  —  The  native  mineral  orpiment 
(As2S3)  is  used  in  making  a  yellow  paint,  and  realgar  (As2S2)  a  red 


ANTIMONY  283 

paint.  Paris  green  (Cu3(AsO3)2.Cu(C2H3O2)2)  and  lead  arsenate 
(Pb3(AsO4)2)  are  effective  insecticides  and  are  used  extensively  to 
exterminate  potato  bugs  'and  other  insect  pests.  Arsine  (AsH3)  is 
a  gas  analogous  to  NH3  and  PH3. 

The  formation  of  yellow  arsenious  sulphide  (As2S3)  by  passing 
hydrogen  sulphide  into  an  arsenic  solution  containing  hydrochloric 
acid  is  the  usual  test  for  arsenic.  (See  Part  II,  Exp.  161.) 

Antimony 

349.  Occurrence.  —  Small  quantities  of  free  anti- 
mony are  found.  The  most  common  native  compound 
is  stibnite  (Sb2S3). 

350.  Antimony  is  prepared  on  a  large  scale  by  two  methods.  In 
one  the  sulphide  is  roasted,  and  the  oxide  thus  formed  is  reduced  with 
charcoal.  Equations  representing  the  main  changes  are  — 


Sb2S3         +         502  Sb204         + 

Antimony  Oxygen  Antimony  Sulphur 

Sulphide  Oxide  Dioxide 

Sb204  +  4C  2Sb  +  4CO 

The  other  method  consists  in  heating  the  sulphide  with  iron,  the 
equation  for  the  chemical  change  being  — 

Sb2S3  +  3Fe       =  2Sb  +          3FeS 

Antimony  Iron  Antimony  Iron 

Sulphide  Sulphide 

351.  Properties.  —  Antimony  is  a  silver-  white,  crystal- 
line, brittle  solid.  Its  specific  gravity  is  6.62.  Antimony 
melts  at  about  630°  C.  At  ordinary  temperatures  it 
does  not  tarnish  in  the  air,  but  when  heated  it  burns 
with  a  bluish  flame,  forming  white,  powdery  antimony 
trioxide  (Sb2O3).  Nitric  acid  oxidizes  it  to  Sb2O3  or  to 
antimonic  acid  (H3Sb04)  and  aqua  regia  transforms  it  into 
soluble  antimony  trichloride  —  the  latter  being  a  conven- 
ient solution  of  antimony  for  use  in  the  laboratory.  (See 
Part  II,  Exps.  165,  166.) 


284  CHEMISTRY 

352.  Alloys  of  Antimony.  —  When  antimony  is  melted 
with  some  metals,  especially  lead  and  tin,  the"  metals  dis- 
solve one  another.     Such  a  metallic  solution  upon  cool- 
ing forms   an   alloy.     Alloys  have   different,   often  very 
different,    properties    from    the    original    metals.     Thus, 
alloys  of  antimony  and  lead  expand  on  cooling  and  are 
used  as  type  metal  because  they  reproduce  sharply  the 
dots  and  fine  lines.     Other  alloys,  like  Babbitt  metal,  are 
used  for  bearings  of  machines. 

353.  Compounds  of  Antimony.  —  Antimony  forms  stibine  (SbH3), 
which  is  analogous  to  ammonia  (NH3).     It  also  forms  complex  com- 
pounds in  which  the  group  SbO  —  called  antimonyl  —  acts  as  a  uni- 
valent   radical,  e.g.  tartar  emetic  or  potassium  antimonyl   tartrate 
(KSbOC4H4O6),  which  is  used  as  a  medicine  and  as  a  mordant  in 
dyeing  cotton.     Antimony  trisulphide  (Sb2S3)  is  obtained  as  an  orange 
red  precipitate  by  passing  hydrogen  sulphide  gas  into  a  solution  of  an 
antimony  salt  —  the  usual  test  for  antimony.     This  sulphide  is  used 
in  making  the  red  rubber  tubing  and  stoppers  used  in  the  laboratory. 
Antimony  trichloride  (SbCl3)  is  formed  by  the  action  of  chlorine  upon 
the  metal  or  by  interaction  with  aqua  regia.     It  hydrolyzes  readily, 
i.e.  interacts  with  water,  thus:  — 

SbCl3          +        H20  SbOCl          +  2HC1 

Antimony  Water  Antimony  Hydrocholoric 

Trichloride  Oxychloride  Acid 

Antimony  oxychloride  is  a  white  solid'  insoluble  in  water,  and  its 
formation  is  sometimes  used  as  a  test  for  antimony.  (See  Part  II, 
Exps.  162,  163.) 

Bismuth 

354.  Occurrence.  —  Bismuth  is  usually  found  in  the 
native  state,  though  it  is  not  abundant  nor  widely  dis- 
tributed.    The   oxide  (bismite,  Bi2O3)  and  the  sulphide 
(bismuthinite,  Bi2S3)  are  the  common  native  compounds. 

355.  Preparation.  —  Bismuth  is  prepared  from  the  native  metal 
by  melting  it  on  an  inclined  plate  and  allowing  it  to  drain  away  from 
the  solid  impurities.     Sometimes  the  sulphide  is  roasted,  and  the 
resulting  oxide  is  reduced  with  charcoal,  as  in  the  case  of  antimony. 


BISMUTH 


285 


356.  Properties.— Bismuth  is  a  silvery  metal  with  a 
reddish  tinge.  Like  antimony,  it  is  very  brittle.  Its 
specific  gravity  is  about  9.9.  It  does  not  tarnish  in  dry 
air,  but  it  grows  dull  in  moist  air;  and  when  heated  in  air 
it  burns  with  a  bluish  flame,  forming  the  yellowish  tri- 
oxide  (Bi2O3).  Hydrochloric  acid  does  not  readily  attack 
it,  but  hot  concentrated  nitric  acid  converts  it  into  a  ni- 
trate and  hot  sulphuric  acid  into  a  sulphate.  Aqua  regia 
transforms  it  into  solu- 
ble bismuth  chloride 
(BiCla).  (See  Part  II, 
Exp.  167.) 

Bismuth  melts  at 
2yi°C.  But  alloys  of 
bismuth,  lead,  and  tin 
melt  at  a  much  lower 
temperature.  For  ex- 
ample, Newton's  metal 
melts  at  94.5°  C.  and 
Rose's  metal  at 93. 8° C.; 
while  Wood's  metal, 
which  contains  the 
metal  cadmium  also, 
melts  at  only  60.5°  C. 
These  metallic  mixtures 
are  called  fusible  met- 
als. They  are  used  in 
making  safety  plugs  for  steam  boilers,  fuses  for  elec- 
trical apparatus,  and  connectors  to  hold  in  place 
automatic  fireproof  doors  and  to  close  temporarily  the 
valves  in  the  automatic  sprinkling  apparatus  frequently 
installed  in  large  buildings  (Fig.  78  —  fusible  metal  at  A). 
(See  Part  II,  Exp.  169.) 


Fig.  78.  —  Sprinkler  Head,  Fusible  Link, 
and  Fire-Proof  Door  Held  in  Place  by 
a  Link  of  Fusible  Metal. 


286  CHEMISTRY 

357.  Compounds  of  Bismuth.  —  Bismuth  trioxide  (Bi2O3)  is  a 
yellowish  crystalline  powder  and  is  used  to  fix  the  gilding  on  porce- 
lain. Bismuth  sulphide  (Bi2S3)  is  obtained  as  a  black  precipitate  by 
passing  hydrogen  sulphide  into  a  solution  of  a  bismuth  salt.  The 
trichloride  (BiCl3)  is  formed  by  the  action  of  chlorine  upon  bismuth, 
or  by  treating  bismuth  with  aqua  regia.  With  an  excess  of  water 
the  trichloride  undergoes  hydrolysis,  forming  basic  bismuth  chloride 
(Bi(OH)2Cl)  which  by  loss  of  water  becomes  bismuth  oxychloride 
(BiOCl).  The  latter  is  a  pearl-white  powder,  insoluble  in  water, 
and  its  formation  is  the  usual  test  for  bismuth.  Bismuth  forms 
the  hydroxides  (Bi(OH)3  and  BiO.OH).  Normal  bismuth  nitrate 
(Bi(N03)3)  when  treated  with  hot  water  forms  basic  bismuth 
nitrate  (Bi(OH)2NO3  or  BiONO3).  The  latter,  often  called  sub- 
nitrate  of  bismuth,  is  a  white  tasteless  powder,  and  is  used  as  a 
medicine  for  dyspepsia  and  as  a  cosmetic.  (See  Part  II,  Exp.  168.) 

EXERCISES 

1.  Discuss  the  occurrence  of  phosphorus. 

2.  Describe  the  manufacture  of  phosphorus  (a)  from  a  phosphate  and 
sulphuric  acid  and  (b)  by  the  electric  method. 

3.  Summarize  the  properties  of  (a)  ordinary  phosphorus  and  (b)  red 
phosphorus.     Why  is  phosphorus  so  named? 

4.  What   is   the   formula  of    (a)  tricalcium   phosphate,    (b)    "  sodium 
phosphate,"    (c)  monohydrogen    calcium    orthophosphate,    (d)  superphos- 
phate of  lime,  (e)  dihydrogen  dicalcium  orthophosphate? 

6.  Discuss  the  relation  of  phosphorus  to  life.  Compare  with  nitrogen 
and  carbon  in  this  respect.  Describe  the  manufacture  of  phosphate  fer- 
tilizer. 

6.  Suggest  an  experiment  to  show  that  bones  contain  calcium  phosphate. 

7.  By  what  other  names  is  arsenic  trioxide  known?     What  is  the  anti- 
dote for  arsenic  poisoning? 

8.  What  is  (a)  Paris  green,  (b)  orpiment,  (c]  realgar,  and  (d)  lead  arse- 
nate?     For  what  is  each  used? 

9.  Describe  a  test  for  (a)  arsenic,  (b)  antimony,  and  (c)  bismuth. 
10.   State  the  uses  of  alloys  of  (a)  antimony  and  (b)  bismuth. 

PROBLEMS 

1.  Calculate  the  weight  of  phosphorus  in  (a)  40  metric  tons  of  calcium 
orthophosphate,  (b)  27  gm.  of  silver  metaphosphate,  and  (c)  2  kg.  of  sodium 
pyrophosphate. 


PROBLEMS  287 

2.  Calculate  the  percentage  composition  of   (a)  arsenic  trioxide   and 
arsenic   pentoxide,   and    (b)  antimony   trichloride   and   antimony     penta- 
chloride.     Show  that  these  two  sets  of  substances  illustrate   the  law  of 
multiple  proportions.     (Use  exact  atomic  weights.) 

3.  How  many  liters  of  air  (containing  21  per  cent  of  oxygen  by  volume) 
will  be  required  to  burn  5  gm.  of  phosphorus?     (Standard  conditions.) 

4.  How  many  gm.  of  phosphorus  can  be  made   by  the  electrothermal 
process  from  a  metric  ton  of  calcium  phosphate  (70  per  cent  pure)? 

5.  Write  the  formula  of  (a)  the  orthophosphate,  (b)  the  metaphosphate, 
and  (c)  the  pyrophosphate  corresponding  to  each  of  the  following  metals: 
Al,  barium,  Cd,  Cu(ic),  lead,  Mg,  silver,  K,  and  Ni.     Calculate  the  weight 
of  phosphorus  in  25  gm.  of  three  ortho-  and  three  metaphosphates. 

6.  Calculate   the  atomic  weight  of  phosphorus,  antimony,  or  bismuth 
from  the  following:   (a)  18.5854  gm.  of  phosphorus  give  42.584  gm.  of  phos- 
phorus pentoxide;    (b)  2.99091  gm.  of  antimony  combine  with  1.9495  gm. 
of  sulphur  and  form  antimony  trisulphide;   and  (c)  16.645  gm-  of  bismuth 
trioxide  give  25.2551  gm.  of  bismuth  sulphate  (Bi2(SO4)3). 


CHAPTER   XXIV 
SODIUM  — POTASSIUM  — AMMONIUM  COMPOUNDS 

Sodium  and  potassium  belong  to  a  natural  family  of 
elements  known  as  the  alkali  metals. 

Sodium 

358.  Occurrence.  —  Sodium  is  not  found  free,  but  its 
compounds  are  abundant  and  widely  distributed,  espe- 
cially sodium  chloride.     Many  rocks,  sea  water  and  min- 
eral waters,  and  salt  deposits  contain  sodium  compounds. 

The  symbol  of  sodium,  Na,  is  from  the  Latin  natrium,  which 
comes  from  the  Greek  natron,  an  old  name  of  sodium  carbonate. 

359.  Preparation.  —  Sodium  is  manufactured  on  a  large 

scale  by  the  electrolysis  of  fused  sodium 
hydroxide.  This  was  the  method  by 
which  the  English  chemist  Davy  iso- 
lated sodium  in  1807.  Figure  79  is  a 
sketch  of  one  form  of  the  apparatus 
now  used  at  Niagara  Falls,  where  many 
electrical  industries  are  located.  The 
body  of  the  steel  cylinder  (S)  rests 
within  a  heated  flue.  The  iron  cath- 

Fig.    79. -Apparatus      °de  (C)  PaSSCS  UP  ^rough  the  bottom 

for  the  Manufacture     of  the  cylinder,  while  several  connected 
of  Sodium  by  the    carbon  or  iron  rods   (AA)   enter  from 

Electrolysis    of    So-      aboye    and    constitute    the    anode       A 
dium  Hydroxide. 

cylindrical  collecting  pot  (P)  terminat- 
ing in  a  wire  gauze  surrounds  the  cathode;  it  prevents  the 


SODIUM  —  POTASSIUM  —  AMMONIUM         289 

electrodes  from  touching,  and  does  not  interfere  with  the 
circulation  of  the  molten  sodium  hydroxide.  The  vessel 
(S)  is  filled  with  sodium  hydroxide,  the  lower  portion  in 
the  neck  (B)  being  solid,  the  upper  part  being  kept  mol- 
ten either  by  heat  genererated  by  the  current  or  by  aux- 
iliary gas  burners.  As  the  electrolysis  proceeds,  sodium 
and  hydrogen  are  liberated  at  the  cathode,  rise,  and  col- 
lect in  P.  The  hydrogen  escapes  to  some  extent  through 
the  cover,  but  enough  always  remains  in  the  upper  part 
of  P  to  protect  the  sodium  from  the  air  while  the  molten 
metal  is  being  removed.  Oxygen  is  liberated  at  the 
anode  and  escapes  through  the  pipe  0  without  coming 
in  contact  with  the  sodium  or  hydrogen. 

360.  Properties.  —  Sodium  is  a  silver- white  metal.  It  is 
so  soft  that  it  can  be  easily  molded  with  the  fingers  and 
cut  with  a  knife.  It  floats  on  water,  since  its  specific 
gravity  is  only  about  0.97.  Heated  in.  air,  it  melts  at 
96°  C.,  and  at  a  higher  temperature  it  volatilizes  and 
burns  with  a  brilliant  yellow  flame,  forming  sodium 
oxide.  This  intense  yellow  color  is  characteristic  of 
sodium  and  its  compounds  and  is  the  usual  test  for  the 
element  (free  or  combined).  In  moist  air  the  bright 
surface  quickly  tarnishes,  and  sodium  as  usually  seen  has 
a  yellow  or  brownish  coating.  It  is,  therefore,  kept  under 
kerosene  or  a  liquid  free  from  water.  (See  Part  II, 
Exps.  12  D,  29,  176.) 

Sodium  decomposes  water  at  ordinary  temperatures,  lib- 
erating hydrogen  and  forming  sodium  hydroxide,  thus :  — 

2Na      +     2H2O  =         2NaOH       +      H2 

Sodium  Water  Sodium  Hydroxide        Hydrogen 

If  held  in  one  place  upon  water  by  filter  paper,  enough 
heat  is  generated  to  set  fire  to  the  hydrogen,  which  burns 


29o  CHEMISTRY 

with  a  yellow  flame,  owing  to  the  presence  of  volatilized 
sodium.  Sodium  like  other  metals  interacts  with  acids, 
forming  salts  and  liberating  hydrogen. 

Sodium  is  used  in  the  manufacture  of  sodium  peroxide  (Na202) 
and  sodium  cyanide  (NaCN). 

361.  Sodium  Chloride,  NaCl,  is  the  most  important 
compound  of  sodium.     It  is  familiar  under  the  name  of 
salt  or  common  salt.     The  presence  of  salt  in  the  ocean, 
in  lakes  and  springs,  and  in  the  earth  is  mentioned  in  the 
oldest  historical  records.     It  is  one  of  the  most  abundant 
substances  and  is  the  chief  source  of  sodium  compounds. 

362.  Preparation  of  Common  Salt.  —  Salt  is  obtained  from  sea 
water,  rock  salt  deposits,  and  brines.     Sea  water  contains  nearly  4 
per  cent  of  salts,  and  three  fourths  of  this  amount  is  sodium  chloride. 
The  water  is  evaporated,  often  by  exposure  to  the  sun,  and  the  salt 
separates  from  the  concentrated  solution.     Deposits  of  salt  are  found 
in  many  parts  of  the  globe,  the  most  important  being  in  England, 
Austria-Hungary,  and  Germany.     In  these  regions  and  some  parts 
of  the  United  States,  the  salt  is  mined  like  other  minerals.     In  the 
United  States  much  salt  is  obtained  from  natural  or  artificial  brines, 
i.e.  from  strong  solutions  of  salt. 

According  to  the  standard  established  by  the  United  States  De- 
partment of  Agriculture,  dry  table  or  dairy  salt  must  not  contain 
over  1.4  per  cent  of  calcium  sulphate,  .5  per  cent  of  calcium  and  mag- 
nesium chlorides,  and  .1  per  cent  of  matter  insoluble  in  water.  The 
dampness  of  salt  is  due  to  traces  of  magnesium  and  calcium  chlorides 
(51).  Pure  salt  does  not  absorb  moisture. 

363.  Properties  and  Uses  of  Common  Salt.  —  Salt  is 
rather  uniformly  soluble  in  water,  100  gm.  of  water  dis- 
solving about  36  gm.  of  salt  at  20°  C.,  and  about  39  gm.  at 
100°  C.  (Fig.  16).     It  crystallizes  in  cubes,  and  does  not 
contain  water  of  crystallization.     These  crystals,  when 
heated,  often  snap  open  sharply  (i.e.  decrepitate),  owing 
to  the  sudden  evaporation  of  inclosed  water.     It  melts  at 


SODIUM  —  POTASSIUM  —  AMMONIUM         291 

about  800°  C.  This  substance  is  an  essential  ingredient 
of  the  food  of  man  and  animals.  Besides  its  domestic 
use,  enormous  quantities  are  used  in  preparing  sodium 
carbonate,  sodium  hydroxide,  and  hydrochloric  acid. 
(See  Part  II,  Exp.  171.) 

364.  Sodium   Carbonate,   Na2CO3,   was    formerly  ob- 
tained from  the  ashes  of  marine  plants,  hence  the  old 
name  soda  ash;  sodium  chloride  is  now  the  source.     The 
manufacture   of   sodium    carbonate   is   a  very  extensive 
chemical  industry. 

365.  The  Leblanc  Process,  which  is  the  older  and  is  used  chiefly 
in  Europe,  involves  three  reactions.     Sodium    chloride  is  changed 
into  sodium  sulphate,  thus:  — 

2NaCl        +        H2S04  Na2SO4        +        2HC1 

Sodium  Sulphuric  Sodium  Hydrochloric 

Chloride  Acid  Sulphate  Acid 

The  sodium  sulphate  is  changed  into  sodium  carbonate  by  heating 
it  with  coal  and  calcium  carbonate;  the  two  main  changes,  which 
are  accomplished  by  one  operation,  are  represented  by  the  equations  — 

Na2SO4         +       2C       =  Na2S  +         2C02 

Sodium  Sulphate  Carbon  Sodium  Sulphide  Carbon  Dioxide 

Na2S  +        CaCOs  Na2CO3         +  CaS 

Sodium  Calcium  Sodium  Calcium 

Sulphide  Carbonate  Carbonate  Sulphide 

The  product  is  a  dark  mass  called  black  ash  from  which  the  sodium 
carbonate  is  rapidly  dissolved.  The  solution  of  sodium  carbonate 
is  evaporated,  and  from  it  separate  crystals  having  the  composition 
Na2CO3.ioH2O  and  known  as  sal  soda  or  soda  crystals.  The  crystals 
are  often  heated  until  the  water  of  crystallization  is  driven  off,  and 
the  product  is  then  called  soda  ash  or  calcined  soda  (Na2CO3). 

366.  The  Solvay  Process,  which  is  the  newer  and  is  operated  very 
successfully  in  the  United  States,  consists  in  saturating  a  cold  con- 
centrated solution  of  sodium  chloride  first  with  ammonia  gas  and  then 
with  carbon  dioxide  gas.     The  equation  for  the  complete  chemical 
change  is  — 

H20   +  NaCl     +     NH3     +     CO2     =     HNaCO3     +     NH4C1 

Water  Sodium  Ammonia  Carbon  Acid  Sodium  Ammonium 

Chloride  Dioxide  Carbonate  Chloride 


292  CHEMISTRY 

The  acid  sodium  carbonate  is  sparingly  soluble  in  cold  ammonium 
chloride  solution,  and  is  therefore  precipitated.  (See  Part  II,  Exp. 
177.)  The  acid  sodium  carbonate  is  changed  into  normal  sodium 
carbonate  by  heating,  thus:  — 

2HNaCO3  Na2C03         +          CO2         +          H20 

Acid    Sodium  Sodium  Carbon  Water 

Carbonate  Carbonate  Dioxide 

367.  Properties   and   Uses   of   Sodium   Carbonate.  - 

Crystallized  sodium  carbonate  (Na2CO3.ioH2O)  is  often 
called  sal  soda  or  washing  soda.  It  slowly  loses  its  water 
of  crystallization  when  exposed  to  air.  When  heated,  it 
dissolves  in  its  water  of  crystallization,  and  continued 
heating  changes  it  into  the  white  anhydrous  salt  (Na2COs). 
It  is  soluble  in  water,  and  the  alkaline  solution  is  widely 
used  as  a  cleansing  agent;  hence  the  name  washing  soda. 
Enormous  quantities  are  used  in  the  manufacture  of  glass, 
soap,  and  many  other  useful  substances  (310,  249). 

The  alkalinity  of  sodium  carbonate  solution  is  due  to  hydrolysis 
(165),  and  is  thus  explained  in  terms  of  the  ionic  hypothesis:  Sodium 
carbonate  ionizes  into  2Na+  and  CO3  ,  but  the  unstable  CO3-ions 
form  HCO3-ions  with  the  H-ions  from  the  slightly  dissociated  water. 
This  removal  of  H-ions  finally  leaves  in  the  solution  sufficient  OH-ions 
to  produce  alkalinity. 

368.  Sodium   Bicarbonate,   HNaC03,   is  prepared  by 
the  Solvay  process  (see  above),  or  by  treating  a  sodium 
carbonate  solution  with  carbon  dioxide  gas.     It  is  some- 
times called  acid  sodium  carbonate  (though  it  is  nearly 
neutral  to  litmus)  and  hydrogen  sodium  carbonate.     It  is 
a  white  powder,  less  soluble  in  water  than  the  normal 
carbonate.     When  heated  alone  or  when  mixed  with  an 
acid  or  an  acid  salt,  sodium  bicarbonate  gives  up  carbon 
dioxide.     This  property  early  led  to  its  use  in  cooking, 
and  gave  the  names  cooking  soda,  baking  soda,  or  simply 
soda. 


SODIUM  —  POTASSIUM  —  AMMONIUM          293 

369.  Baking  Powder.  —  Sodium  bicarbonate  is  an  essential  ingre- 
dient of  baking  powder  and  of  the  various  mixtures  (except  yeast) 
used  to  raise  bread,  cake,  and  other  food.     The  other  ingredient  is 
usually  a  mild  acid  salt,  such  as  an  acid  phosphate  (340)  or  cream  of 
tartar  (acid  potassium  tartrate  (HKC4H4O6)  ),  which  slowly  liberates 
the  carbon  dioxide  from  the  sodium  bicarbonate  (243).     Sour  milk, 
which  contains  lactic  acid,  is  sometimes  used  in  place  of  cream  of 
tartar  (241).     When  pastry  is  raised  with  baking  powder  or  a  mixture 
of  baking  soda  and  cream  of  tartar,  the  escaping  carbon  dioxide  puffs 
up  the  unbaked  mixture.     Hence  baking  soda  is  often  called  saleratus 

—  the  salt  that  aerates  (from  the  Latin  words  sal,  salt,  and  aer,  air  or 
gas).  (See  Part  II,  Exp.  113.) 

370.  Sodium  Hydroxide  or  Caustic  Soda,  NaOH,  is 
a  white,  crystalline,  brittle,  corrosive  solid.     It  absorbs 
water  and  carbon  dioxide  rapidly  from  the  air,  and  is 
thereby   changed   into   sodium   carbonate.     It   dissolves 
readily  in  water,  with  rise  of  temperature.     The  solution 
is  strongly  alkaline  and  disintegrates  many  substances, 
hence  the  term  caustic.     It  melts  easily.     Immense  quan- 
tities are  used  in  making  hard  soap,  paper,  and  dyestuffs. 

371.  Manufacture  of  Sodium  Hydroxide.  —  The  chemical  process 

consists  in  boiling  a  dilute  solution  of  sodium  carbonate  with  calcium 
hydroxide;  the  main  change  is  represented  thus:  — 

Ca(OH)2     +     Na2CO3     =     2NaOH     +     CaCO3 

Calcium  Sodium  Sodium  Calcium 

Hydroxide  Carbonate  Hydroxide  Carbonate 

The  solution  of  sodium  hydroxide  is  evaporated  and  the  molten  mass 
is  allowed  to  solidify  in  small  cylindrical  molds  about  the  diameter 
of  a  lead  pencil  or  in  large  iron  barrels  called  drums.  In  the  elec- 
trolytic process,  which  is  operated  on  a  large  scale  at  Niagara  Falls, 
New  York,  a  solution  of  sodium  chloride  is  used.  One  form  of  appa- 
ratus is  shown  in  Fig.  80.  It  consists  of  a  slate  box  divided  into  one 
cathode  and  two  anode  compartments  by  partitions  extending  nearly 
to  the  bottom;  the  compartments  are  separated  by  a  layer  of  mercury 
(shown  in  black).  The  T-shaped  anodes  (A,  A)  of  graphite  and  the 
cathode  (C)  of  iron  reach  nearly  to  the  mercury.  The  anode  com- 


294 


CHEMISTRY 


partments  contain  sodium  chloride  solution,  while  the  cathode  com- 
partment contains  sodium  hydroxide  solution;  sodium  chloride 
solution  of  the  right  concentration  flows  slowly  and  continuously 
through  the  anode  compartments  by  means  of  the  pipes  E,  E  (and 
outlets  not  shown).  When  the  current  passes,  chlorine  is  evolved  at 


Fig.  80.  —  Apparatus  for  the  Manufacture  of  Sodium 
Hydroxide  by  Electrolysis  of  Sodium  Chloride. 

the  anodes  and  escapes  through  the  pipes  D,  D;  the  sodium  is  liber- 
ated at  the  intermediate  cathode  of  mercury  and  forms  an  amalgam 
with  it.  By  carefully  rocking  the  cell  on  the  device  X,  X  the  sodium 
amalgam  alone  flows  beneath  the  partitions  into  the  cathode  com- 
partment, where  the  sodium  is  liberated  at  the  iron  cathode;  the 
sodium  at  once  reacts  with  the  water  forming  hydrogen  and  sodium 
hydroxide.  The  hydrogen  escapes  through  the  pipe  H,  while  the 
sodium  hydroxide  solution  is  allowed  to  become  concentrated,  being 
then  drawn  off  (through  G)  and  replaced  by  water  (through  F). 
The  sodium  hydroxide  solution  is  treated  as  described  above,  while 
the  chlorine  is  stored  in  steel  cylinders  or  used  directly  in  manufactur- 
ing bleaching  powder  and  other  chlorine  compounds  (75,  79,  80). 

372.  Sodium  Sulphate,  Na2SO4,  is  a  white  solid.  It 
dissolves  readily  in  water,  and  when  a  strong  solution 
made  at  30°  C.  is  cooled,  large  transparent,  bitter  crystals 
separate.  They  have  the  formula  Na2SO4.ioH2O  and 
are  called  Glauber's  salt.  Large  quantities  are  used  in 
making  sodium  carbonate  and  glass  (365,  309,  310). 

Sodium  sulphate  is  prepared  by  the  interaction  of  sodium  chloride 
and  sulphuric  acid  (83,  365)  or  magnesium  sulphate.  The  equation 
for  the  latter  reaction  is  — 


SODIUM  —  POTASSIUM  —  AMMONIUM         295 

MgS04        +       2NaCl         =       Na2S04        +       MgCl2 

Magnesium  Sodium  Sodium  Magnesium 

Sulphate  Chloride  Sulphate  Chloride 

373.  Sodium  Nitrate,  NaNO3,  is  found  abundantly  in 
Chile  and  is  often  called  Chile  saltpeter.     It  is  a  white, 
or  brownish,  solid,  which  becomes  moist  in  the  air.     Large 
quantities  are  used  as  a  fertilizer,  either  alone  or  mixed 
with  compounds  of  potassium  and  of  phosphorus  (98,  344, 
386,  and  compare  106).     It  is  used  in  making  nitric  and 
sulphuric  acids,  and  potassium  nitrate. 

The  deposits  of  sodium  nitrate  are  in  a  dry  region  near  the  coast 
and  cover  a  large  area.  Chile  controls  the  industry  and  exports 
annually  over  a  million  tons.  The  crude  salt,  which  is  called  caliche, 
looks  like  rock  salt.  The  commercial  salt  is  extracted  from  caliche 
by  treating  with  water,  settling,  and  evaporating  the  solution  of  the 
nitrate  to  crystallization.  The  final  mother  liquor  is  a  source  of 
iodine  (327). 

374.  Sodium  Dioxide  or  Peroxide,  Na2O2,  is  a  yellowish  solid. 
It  is  used  to  bleach  straw  and  delicate  fabrics.     With  water  it  liber- 
ates oxygen,  according  to  the  equation  — 

Na202          +      H20     =        O         +          2NaOH 

Sodium  Dioxide  Water  Oxygen  Sodium  Hydroxide 

375.  Other    Sodium    Compounds.  —  Sodium    phosphates    (340), 
sodium  thiosulphate  (292),  acid  sodium  sulphite  (280,  281),  sodium 
silicate   (306),   and  sodium  tetraborate  or  borax   (295)   have  been 
described. 

Potassium 

376.  Occurrence.  —  This  metal  is  not  found  free,  but 
its  compounds  are  abundant.     The  minerals  mica  and 
feldspar  are  silicates  containing  potassium.     By  the  decay 
of  these  and  other  minerals,  potassium  compounds  find 
their  way  into  the  soil  and  are  taken  up  by  plants.     Potas- 
sium salts  are  found  in  wood  ashes  and  in  the  deposits 
in  wine  casks.     Sea  water  and  mineral  waters  contain 
potassium    salts,    particularly    potassium    chloride    and 


296  CHEMISTRY 

potassium  sulphate.     Extensive  beds  of  potassium  salts 
are  found  in  Germany,  especially  at  Stassfurt. 

377.  The  Stassfurt  deposits  of  the  salts  of  potassium  and  other 
metals  consist  of  about  thirty  different  salts.  The  deposits  were 
doubtless  formed  ages  ago  by  the  evaporation  of  an  inclosed  arm  of 
the  sea  under  special  conditions.  The  lowest  bed  is  an  enormous 
mass  of  rock  salt.  Upon  this  rest  more  or  less  regular  layers  of 
potassium  and  magnesium  salts,  and  higher  still  are  calcium  salts. 
The  most  important  Stassfurt  potassium  minerals  are:  — 

Sylvite  —  KC1  Carnallite  —  KC1,  MgCl2.6H20 

Kainite  —  KC1,  MgSO4.3H2O        Picromerite  —  K2SO4,  MgSO4.6H2O 

378.  Preparation  and  Properties.  —  Potassium  is  pre- 
pared from  potassium  hydroxide  by  electrolysis.  It  was 
first  obtained  by  this  method  in  1807  by  Davy.  Potas- 
sium is  a  soft,  silver-white  metal  with  a  slight  bluish 
tinge.  It  floats  on  water,  the  specific  gravity  being  about 
0.86.  Its  brilliant  luster  soon  disappears  in  air,  owing  to 
rapid  oxidation.  Potassium  as  ordinarily  seen  is,  there- 
fore, covered  with  a  grayish  coating,  and,  like  sodium, 
must  be  kept  under  mineral  oil.  It  melts  at  62.5°  C.,  and 
at  a  higher  temperature  burns  with  a  violet-colored  flame. 
This  color  is  characteristic  of  burning  potassium,  and  is 
a  test  for  the  metal  and  its  compounds.  It  interacts 
with  water  more  energetically  than  sodium  (see  Fig.  5). 

379.  Potassium  Chloride,  KC1,  is  a  white  solid  which 
crystallizes   in    cubes    and    otherwise    resembles    sodium 
chloride.     It  is  found  native  in  the  Stassfurt   deposits. 
Considerable  is  obtained  by  decomposing  carnallite  and 
crystallizing  the  potassium  chloride  from  the  solution  of 
the  more  soluble  magnesium  chloride.     It  is  used  chiefly 
to  prepare  other  potassium  salts  and  as  a  fertilizer  (380, 
383,  386). 

380.  Potassium   Nitrate,   KNO3,  is   also  called  niter 


SODIUM  —  POTASSIUM  —  AMMONIUM         297 

and  saltpeter.  It  is  formed  in  the  soil  of  many  warm 
countries  by  the  decomposition  of  nitrogenous  organic 
matter  (112). 

Potassium  nitrate  is  a  white  solid.  It  dissolves  easily 
in  cold  water  with  a  marked  fall  of  temperature,  and 
very  freely  in  hot  water  (Fig.  16).  Unlike  sodium  nitrate 
it  does  not  become  moist  in  the  air.  It  crystallizes 
readily,  but  contains  no  water  of  crystallization.  The 
taste  is  salty  and  cooling.  It  melts  at  333°  C.,  and  on 
further  heating  changes  into  potassium  nitrite  (KNO2) 
and  oxygen.  At  a  high  temperature,  potassium  nitrate 
gives  up  oxygen  readily,  especially  to  charcoal,  sulphur, 
and  organic  matter.  This  oxidizing  power  leads  to  its 
extensive  use  in  making  gunpowder,  fireworks,  matches, 
explosives,  and  in  many  chemical  operations.  It  is  also 
used  in  salting  meat. 

Potassium  nitrate  is  manufactured  by  mixing  hot,  concentrated 
solutions  of  sodium  nitrate  and  potassium  chloride.  The  equation 
for  the  reaction  is  — 

NaN03     +       KC1         =       KN03      +      NaCl 

Sodium  Potassium  Potassium  Sodium 

Nitrate  Chloride  Nitrate  Chloride 

The  sodium  chloride,  being  much  less  soluble  than  the  potassium 
nitrate,  separates  and  is  removed  by  filtration.  By  evaporating  the 
filtrate,  crystals  of  potassium  nitrate  separate  and  are  further  puri- 
fied by  fecrystallization.  The  solubility  of  these  salts  is  shown  in 
Fig.  16.  (See  Part  II,  Exp.  179.) 

381.  Gunpowder  is  a  mixture  of  potassium  nitrate,  charcoal,  and 
sulphur.  The  proportions  differ  with  the  use  of  the  powder.  A 
common  variety  contains  75  per  cent  of  potassium  nitrate,  15  of  char- 
coal, and  10  of  sulphur.  When  gunpowder  burns  in  a  closed  space, 
hot  gases  are  suddenly  formed.  The  pressure  exerted  by  these  gases 
forces  the  bullet  from  a  gun  and  tears  rocks  to  pieces.  The  chemical 
changes  attending  the  explosion  of  gunpowder  in  a  closed  space  are 
complex,  as  may  be  seen  by  the  following  (approximate)  equation:  — 


298  CHEMISTRY 

i6KNO3     +     2iC     +     58  =  5K2CO3     +     i3C02     +     K2SO4 

Potassium  Carbon  Sulphur     Potassium  Carbon  Potassium 

Nitrate  Carbonate  Dioxide  Sulphate 

+  3CO     +      2K2S2     +     8N2 

Carbon  Potassium  Nitrogen 

Monoxide  Bisulphide 

The  equation  for  the  explosion  when  unconfined  is  much  simpler, 
thus:  - 

2KN03  +  3C  +  S  =  3C02  +  N2  +     K2S 

Potassium 
Monosulphide 

Gunpowder  is  being  rapidly  replaced  by  smokeless  powders  (230). 

382.  Potassium  Chlorate,  KC1O3,  is  a  white,  crystal- 
line, lustrous  solid.  It  melts  at  370°  C.,  and  at  a  high 
temperature  decomposes  into  oxygen  and  potassium 
chloride  as  final  products.  It  is  used  to  prepare  oxygen, 
and  in  the  manufacture  of  matches  and  fireworks.  In 
the  form  of  " chlorate  of  potash  tablets"  it  is  used  as  a 
remedy  for  sore  throat. 

Potassium  chlorate  is  manufactured  by  the  electrolysis  of  potas- 
sium chloride  in  a  special  apparatus  which  allows  the  two  products  — 
chlorine  and  potassium  hydroxide  —  to  mix  and  interact.  The 
equation  for  the  reaction  is  — 


3C12     +     6KOH     =     KC103     +     5KC1     + 

Chlorine  Potassium  Potassium  Potassium  Water 

Hydroxide  Chlorate  Chloride 

Another  process  consists  in  preparing  calcium  chlorate  and  convert- 
ing this  into  potassium  chlorate  by  interaction  with  potassium  chlo- 
ride. The  equations  for  the  changes  are  — 


6Ca(OH)2     +     6C12     =     Ca(ClO3)2     +     5CaCl2     +     6H20 

Calcium  Calcium  Calcium 

Hydroxide  Chlorate  Chloride 

Ca(ClO3)2  +  2KC1  =  2KC103  +  CaCl2 

383.  Potassium  Carbonate,  K2CO3,  is  a  white  solid. 
It  deliquesces  in  the  air,  is  very  soluble  in  water,  and  the 
solution,  like  that  of  sodium  carbonate,  has  a  strong  alka- 
line reaction  (165,  367)  .  It  was  formerly  obtained  by  treat- 


SODIUM  —  POTASSIUM  —  AMMONIUM         299 

ing  wood  ashes  with  water  and  evaporating  the  solution 
to  dryness.  The  crude  salt  thus  obtained  has  long  been 
called  potash,  and  a  purer  product  is  known  as  pearlash. 
The  name  of  the  element  potassium  was  suggested  by  the 
word  potash;  the  symbol,  K,  comes  from  the  Latin 
word  kalium.  It  is  used  extensively  in  the  manufacture 
of  hard  glass,  soft  soap,  caustic  potash  (potassium  hy- 
droxide), and  other  potassium  compounds.  Potassium 
carbonate  is  made  from  potassium  chloride  by  the  Le- 
blanc  or  other  processes. 

384.  Potassium  Hydroxide  or  Caustic  Potash,  KOH, 
is  a  white,  brittle  solid,   resembling  sodium  hydroxide. 
It  absorbs  water  and  carbon  dioxide  very  readily;   and  if 
exposed  to  the  air,  soon  becomes  a  thick  solution  of  potas- 
sium  carbonate.      Like   sodium   hydroxide,   it   dissolves 
readily  in  water  with  evolution  of  heat,  forming  a  strongly 
alkaline  solution.     Its  solutions  corrode  and  disintegrate 
animal    and  vegetable    matter  and  many  mineral  sub- 
stances   such    as    glass   and  porcelain;    hence   the   term 
caustic  potash.     Besides  its  use  in  the  laboratory,  large 
quantities  are  consumed  in  making  soft  soap.     It  is  pre- 
pared from  potassium  chloride  by  the  methods  used  for 
sodium  hydroxide. 

385.  Other  Potassium  Compounds.  —  Potassium  Cya- 
nide (KCN)  is  a  white  solid,  very  poisonous,  very  soluble 
in  water,  and  having  an  odor  like  bitter  almonds.     Potas- 
sium Sulphate  (K2SO4),  which  is  a  white  solid  resembling 
sodium  sulphate,  is  manufactured  from  kainite  and  other 
Stassfurt  salts.     It  is  largely  used  as  a  fertilizer  and  in 
making  potassium  carbonate  and  potassium  alum  (437). 

386.  Relation  of  Potassium  to  Life.  —  Potassium,  like 
nitrogen  and  phosphorus,  is  essential  to  the  life  of  plants 
and  animals.     The  ash  of  many  common  grains,  vege- 


300  CHEMISTRY 

tables,  and  fruit  contains  potassium  carbonate,  which  is 
formed  from  complex  organic  potassium  compounds. 
Potassium  salts  taken  from  the  soil  by  plants  must  be 
returned  if  the  soil  is  to  be  productive.  Sometimes  crude 
kainite  is  used  extensively  as  a  fertilizer  (373) ;  wood  ashes, 
or  the  sulphate  and  chloride,  are  often  used  to  supply 
potassium  salts. 

387.  The  Alkali  Metals.     The  properties  of  the  metals  are  quite 
similar,   and  the  chemical  activity  increases  from  lithium   (at.  wt. 
6.94)   to  caesium   (at.  wt.   132.81).     All  decompose  water,  yielding 
hydrogen  and  an  hydroxide.     The  hydroxides  are  active  bases,  and 
the  familiar  ones  long  ago  gave  the  name  alkali  to  the  family.     Anal- 
ogous compounds  are  much  alike. 

Ammonium  Compounds 

388.  Introduction.  —  Ammonium  (NH4)  is  a  metallic 
radical,  i.e.  a  group  of  elements  which  acts  like  an  atom 
of  a  metal  in  chemical  changes  (105).     Its  most  familiar 
compound  is  ammonium  hydroxide  (NH4OH),  which  has 
the  properties  of  a  base  and  resembles  sodium  and  potas- 
sium hydroxides  (99-105).     Other  compounds  of  ammo- 
nium are  analogous  to  the  corresponding  salts  of  sodium 
and  potassium.     Ammonium  salts  volatilize  when  heated. 
They  also  give  off  ammonia  gas  when  heated  with  an 
alkali,  such  as  calcium  or  sodium  hydroxide  (99).     This 
reaction  is  a  test  for  ammonium  compounds. 

389.  Ammonium    Chloride,    NH4C1,    is    prepared    by 
passing  ammonia  gas  into  dilute  hydrochloric  acid,  by 
mixing  ammonium  hydroxide  and  hydrochloric  acid,  or 
by  letting  ammonia  and  hydrogen  chloride  mingle  (86). 
The  equation  for  the  essential  reaction  is  - 

NH3      +          HC1  NH4C1 

Ammonia        Hydrogen  Chloride         Ammonium  Chloride 


SODIUM  —  POTASSIUM  —  AMMONIUM         301 

It  is  convenient  to  regard  this  compound  as  the  ammonium 
salt  of  hydrochloric  acid,  as  if  it  were  formed  by  replacing 
the  hydrogen  of  the  acid  by  ammonium,  just  as  sodium 
forms  sodium  chloride. 

Ammonium  chloride  is  a  white,  granular,  fibrous,  or 
crystalline  solid,  with  a  sharp,  salty  taste.  It  dissolves 
easily  in  water,  and  in  so  doing  lowers  the  temperature 
markedly.  (See  Part  II,  Exp.  175.) 

Large  quantities  of  ammonium  chloride  are  made  by  liberating 
ammonia  from  the  ammoniacal  liquor  obtained  in  gas  works  and  pass- 
ing the  gas  into  hydrochloric  acid  (205).  The  crude  product  is  often 
called  "muriate  of  ammonia"  to  indicate  its  relation  to  muriatic  (or 
hydrochloric)  acid.  It  is  largely  used  in  charging  Leclanche  bat- 
teries, as  an  ingredient  of  soldering  fluids,  in  galvanizing  iron,  and  in 
textile  industries.  The  crude  salt  is  purified  by  heating  it  gently  in 
a  large  iron  or  earthenware  pot,  with  a  dome-shaped  cover;  the 
ammonium  chloride  volatilizes  easily  and  then  crystallizes  in  the 
pure  state  as  <J  fibrous  mass  on  the  inside  of  the  cover,  but  the  im- 
purities remain  behind  in  the  vessel.  This  process  of  purification 
is  called  sublimation  (328).  The  product  is  called  a  sublimate. 
Sublimed  ammonium  chloride  is  known  as  sal  ammoniac. 

390.  Ammonium    Sulphate,    (NH4)2SO4,   is   made   by 
passing  ammonia  gas  into  sulphuric  acid,  or  by  adding 
ammonium  hydroxide  to  the  acid,  thus:  - 

2NH4OH      +     H2SO4    =       (NH4)2SO4      +     2H2O 

Ammonium  Hydroxide  Ammonium  Sulphate 

The  commercial  salt  is  a  grayish  or  yellowish  solid  and  is 
obtained  in  large  quantities  as  a  by-product  of  coal  gas 
manufacture  (205).  It  is  used  as  an  ingredient  of  fer- 
tilizers, since  it  is  an  inexpensive  salt  containing  consid- 
erable nitrogen. 

* 

391.  Ammonium  Nitrate,  NH4NO3,  is  made  by  passing  ammonia 
into  nitric  acid,  or  by  allowing  ammonia  gas  and  the  vapor  of  nitric 
acid  to  mingle,  thus:  — 


302  CHEMISTRY 

NH3        +       HN03       =          NH4N03 

Ammonia  Nitric  Acid  Ammonium  Nitrate 

It  is  a  white  salt  which  forms  beautiful  crystals.  It  dissolves  easily 
in  water  with  a  fall  of  temperature.  When  gently  heated  it  decom- 
poses into  nitrous  oxide  (N2O)  and  water,  and  its  chief  use  is  in  the 
preparation  of  nitrous  oxide  (115).  (See  Part  II,  Exp.  56.) 

392.  Ammonium  Carbonate,  (NH4)2CO3,  is  a  white  transparent 
solid  when  pure  and  fresh,  but  on  exposure  to  air  it  loses  ammonia 
and  becomes  opaque.  It  is  used  in  some  kinds  of  baking  powder, 
to  scour  wool,  as  a  medicine,  and  to  prepare  smelling  salts  (since  it 
gives  off  ammonia  readily). 

EXERCISES 

1.  Describe  the  manufacture  of  sodium  by  electrolysis. 

2.  Summarize  the  physical  properties  and  chemical  properties  of  sodium. 

3.  Discuss    the  manufacture  of  .  sodium  carbonate  by  (a)  the  Leblanc 
process,     (b)  By  the  Solvay  process. 

4.  What  is  (a)  soda,  (b)  soda  ash,  (c)  sodium  carbonate,  (d)  soda  crys- 
tals, (e)  sal  soda,  (/)  washing  soda,  (g)  "alkali,"  (//)  acid  sodium  carbonate, 
(»')  saleratus,  (j)  baking  powder,  (k)  baking  soda,  (/)  caustic  soda? 

5.  Describe  the  manufacture  of  sodium  hydroxide. 

6.  What  is  a  simple  test  for  combined  (a)  sodium,  and  (b)  potassium? 

7.  Describe  the  preparation  of  potassium  nitrate. 

8.  Suggest  an  experiment  to  find  the  per  cent  of  ammonium  chloride 
in  a  mixture  of  sodium  and  ammonium  chlorides. 

PROBLEMS 

1.  Calculate  the  weight  of  sodium  in  (a)  20  gm.  of  NaOH,  (b)  35  gm. 
of  Na2CO3,  (c)  22  gm.  of  acid  sodium  sulphate,  and  (d)  18  gm.  of  sodium 
nitrate. 

2.  Calculate  the  weight  of  potassium  in  (a)  15  gm.  of  KC1,  (b)  37  gm. 
of  potassium  carbonate,  (c)  24  gm.  of  KNO3,  and  (d)  45  gm.  of  potassium 
hydroxide. 

3.  Find  the  weight  of  ammonium  in  (a)  20  gm.  of  ammonium  hydroxide, 
(b)  50  gm.  of  ammonium  nitrate,  (c)  60  gm.  of  ammonium  sulphate,  and 
(d)  170  gm.  of  sal  ammoniac. 

4.  Write  the  formulas  of  the  sodium,  potassium,  and  ammonium  salts 
of  the  following  acids:  chloric,  phosphoric (ortho),  sulphurous,  hydrofluoric, 
carbonic,    acetic,   hydrobromic,    dichromic,   permanganic,    manganic. 

5.  Indicate  the  ions  in  dilute  aqueous  solutions  of  (a)  sodium  hydroxide, 
(b)  potassium  nitrate,    (c)  sodium   chloride,   and  (d)  ammonium   chloride. 


CHAPTER  XXV 
COPPER  —  SILVER  —  GOLD 

393.  Introduction.  —  Copper,    silver,    and    gold    have 
been  known  for  ages.     Domestic  utensils  and  weapons 
containing  copper  were  used   before   similar  objects   of 
iron.     Silver  and  gold  have  always  been  regarded  as  pre- 
cious metals  and  have  long  been  used  for  ornaments  and 
money.     The  Latin  words  cuprum,  argentum,  and  aurum 
give  us  the  symbols  Cu,  Ag,  and  Au. 

394.  Occurrence  of  Copper.  —  Copper,  both  free  and 
combined,  is  an  abundant  element.     Like  many  other 
metals  copper  occurs  as  ores,  that  is,  mineral  compounds 
from  which  the  metal  can  be  profitably  extracted;   some- 
times  the   term  ore  is  applied  to  rock  containing  metal, 
e.g.  quartz-bearing  gold  is  called  gold  ore.     Free  or  native 
copper  is  mined  in  large  quantities  in  northern  Michigan 
on  the  shores  of  Lake  Superior.     The  important  copper 
ores  are  copper  sulphide  (chalcocite,  copper  glance,  Cu2S), 
copper  oxide  (cuprite,  ruby  ore,  Cu2O),  the  copper-iron 
sulphides    (copper    pyrites,    chalcopyrite,    CuFeS2,    and 
bornite,  Cu3FeS3),  and  the  complex  carbonates  (malachite, 
CuCO3.Cu(OH)2,  and  azurite,  2CuCO3.Cu(OH)2).     In  the 
United  States  copper  is  produced  chiefly  in  Montana, 
Michigan,  and  Arizona,  the  annual  output   being   over 
half  the  world's  supply. 

395.  Metallurgy  of  Copper.  —  Metallurgy  is  the  art  of 
extracting  metals  from  their  ores.     And  since  the  pro- 
cesses  applied   to   copper   ores   involve   several   general 
operations,  the  extraction  of  copper  is  a  good  example  of 


304  CHEMISTRY 

metallurgy.  First,  the  valuable  part  of  the  ore  is  sepa- 
rated from  the  earthy  or  rocky  impurities  by  crushing, 
washing,  and  mechanical  separation;  this  preliminary 
treatment  is  called  concentration  and  is  often  supple- 
mented by  operations  which  prepare  the  ore  for  chemical 
treatment.  The  next  general  step  is  the  extraction  of  the 
metal  in  a  more  or  less  pure  form  from  the  concentrated 
and  prepared  ore.  This  operation  usually  consists  in 
roasting  the  ore,  and  then,  if  necessary,  heating  the 
product  with  carbon  (or  some  other  reducing  agent)  and 
a  flux,  e.g.  sand  or  limestone,  which  interacts  with  impuri- 
ties and  forms  a  fusible  substance  called  slag,  which  can 
be  easily  separated  and  removed.  The  extraction  of  the 
metal  by  heating  the  ore  is  called  smelting.  Sometimes 
the  metal  is  removed  by  dissolving  agents.  Other  opera- 
tions are  often  necessary  and  they  vary  with  the  ore  and 
the  metal. 

Copper  carbonates  and  oxides  are  reduced  by  heating 
them  with  coke  in  a  suitable  furnace.  The  general 
chemical  change  may  be  represented  thus :  — 

Cu2O       +      C      =    2Cu   +  CO 

Copper  Oxide          Carbon        Copper        Carbon  Monoxide 

The  smelting  of  copper-iron  sulphides  is  complicated. 
After  concentration  and  roasting,  the  mass  is  smelted  with 
coal  and  sand  in  a  furnace.  This  operation  converts 
much  of  the  sulphur  into  sulphur  dioxide  and  most  of  the 
iron  into  a  fusible  iron  silicate,  which  is  removed  as  slag. 
The  remaining  matte,  as  it  is  called,  which  contains  from 
35  to  50  per  cent  of  copper,  besides  iron,  sulphur,  silver, 
and  gold,  must  be  further  treated  to  obtain  the  copper. 
In  some  cases  it  is  roasted  and  melted  until  all  the  iron 
is  removed  and  mainly  copper  sulphide  remains;  this 


COPPER,—  SILVER  —  GOLD 


305 


mass  is  finally  roasted  to  convert  it  partly  into  an  oxide, 
and  the  mixture  of  sulphide  and  oxide  is  again  melted; 
the  sulphur  passes  off  as  sulphur  dioxide,  and  the  copper 
is  left  behind.  The  equation  for  this  final  change  is  - 


2CuO 

Copper  Oxide 


Cu2S          =    4Cu 

Copper  Sulphide         Copper 


SO2 

Sulphur  Dioxide 


Usually  the  matte  is  melted  and  poured  into  a  barrel- 
shaped  vessel,  called  a  converter,  lined  with  silica  and  pro- 
vided with  holes  near  the  bottom.  Air  is  blown  through 
the  liquid,  thereby  removing  the  remaining  sulphur  and 
iron  as  sulphur  dioxide  and  slag  and  leaving  the  copper 
nearly  pure.  The  metal,  which  is  about  98  per  cent  pure, 
is  poured  into  anode  molds,  and  subsequently  sent  to  the 
refinery. 

396.  Refining  of  Copper  by  Electrolysis.  —  Since  very 
pure  copper  is  needed  for  many  purposes,  especially  in 
electrical  industries,  the 
blister  copper,  as  it  is 
called,  which  is  prepared 
by  the  process  described 
above,  must  be  further 
purified.  This  is  done  by 
electrolysis,  and  the  re- 
fined metal,  which  is  ex- 


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Fig.  81.  —  Apparatus  for  the  Prepara- 
tion of  Pure  Copper  by  Electrolysis. 
A,  A,  A  are  Anodes,  and  C,  C,  C  are 
Cathodes. 


ceedingly  pure,  is  called 
electrolytic  copper. 

The  anodes  are  suspended  in  a  solution  of  copper  sul- 
phate. Sheets  of  pure  copper  are  made  cathodes  and  dip 
into  the  solution,  as  shown  in  Fig.  81.  When  the  current 
passes,  pure  copper  dissolves  from  the  anodes  and  deposits 
upon  the  cathodes,  the  impurities  either  remaining  in 
solution  or  falling  to  the  bottom  of  the  cell  as  a  slime; 


3o6  CHEMISTRY 

from  this  slime  the  gold  and  silver  that  were  in  the  copper 
ore  are  extracted.  In  terms  of  the  theory  of  electrolytic 
dissociation,  copper  ions  (Cu++)  migrate  to  the  cathode, 
lose  their  charges,  and  become  metallic  copper  (166). 

397.  Properties  of  Copper.  —  Copper  is  distinguishable 
from  all  other  metals  by  its  peculiar  reddish  color.     It 
is  flexible,  ductile,  malleable,  and  tough.     It  melts   at 
1083°  C.    Its  specific  gravity  is  8.9.    Copper  is  an  excel- 
lent conductor  of  electricity,  the  pure  metal  being  next 
to  silver  in  this  respect.     Exposed  to  dry  air,  it  turns  dull, 
and  in  moist  air  it  gradually  becomes  coated  with  a  green- 
ish copper  carbonate.     Heated  in  the  air,  it  is  changed 
into  black  copper  oxide  (CuO),  and  at  a  high  tempera- 
ture it  colors  a  flame  emerald-green.     With  nitric  acid  it 
forms  copper  nitrate  and  oxides  of  nitrogen  (112);    with 
hot  sulphuric  acid  it  yields  copper  sulphate  and  sulphur 
dioxide  (277).     Hydrochloric  acid  has  little  effect  upon 
it.     Copper  displaces  some  metals  if  suspended  in  solu- 
tions of  their  compounds,  e.g.   a  clean  copper  wire,  if 
placed  in   a  solution   of  any  mercury  compound,   soon 
becomes  coated  with  mercury;   on  the  other  hand,  metals 
like  iron,  zinc,  and  magnesium  -displace  copper  from  its 
solution,  e.g.  a  nail  or  knife  blade  soon  becomes  coated 
with  copper  if  dipped  into  a  solution  of  any  copper  com- 
pound.    Scrap  iron  is  often  used  to  precipitate  copper 
on  a  large  scale.     (See  414.)     (See  Part  II,  Exps.  181, 184.) 

398.  Tests  for  Copper.  —  The  reddish  color,  peculiar  "coppery" 
taste,  and  green  color  imparted  to  a  flame  serve  to  identify  metallic 
copper.     An  excess  of  ammonium  hydroxide  added  to  a  solution  of 
a  copper  compound  produces  a  beautiful  deep  blue  solution.     A  few 
drops  of  acetic  acid  and  potassium  ferrocyanide  solution  added  to 
a  dilute  solution  of  a  copper  compound  produce  a  brown  precipitate 
of  copper  ferrocyanide  (Cu2Fe(CN6).     (See  Part  II,  Exp.  182.) 


COPPER  —  SILVER  —  GOLD 


307 


399.  Uses  of  Copper.  —  Large  quantities  of  wire  are 
used  in  operating  the  telegraph,  cable,  telephone,  electric 
railway,  and  electric  light.  Sheet  copper  is  made  into 
household  utensils,  boilers,  and  stills,  and  is  also  used  for 
roofing  and  sheathing.  All  nations  use  copper  as  the 
chief  ingredient  of  small  coins.  Much  copper  is  utilized  in 
electrical  and  other  apparatus,  especially  now  that  copper 
can  be  cast  (294).  Books  are  printed  and  illustrated 
from  electrotype  plates,  made  by  depositing  copper  upon 
an  impression  of  the  type  or  design  in  wax;  in  a  similar 
way  many  objects  are  copper  plated  (167).  Copper  is  an 
ingredient  of  many  common  and  useful  alloys,  as  may  be 
seen  from  the  following:  - 

TABLE  or  COPPER  ALLOYS 


Name 

Copper 

Zinc 

Alumin- 

Silver 

Nickel 

Tin 

ium 

Aluminium  bronze    .  . 

90-98 

2-10 

Brass 

63—73 

27—37 

Bronze    

70-95 

1-25 

1-18 

German  silver 

CQ—  6O 

20 

20-25 

Gun  metal    

90 

10 

Gold  coin 

IO 

[Gold  90] 

Nickel  coin  

75 

25 

Silver  coin    

10 

90 

Speculum  metal  

70 

30 

400.  Compounds  of  Copper.  —  Copper  forms  two  series 
of  compounds  —  the  cuprous  and  the  cupric,  the  latter 
being  more  common  under  ordinary  conditions.  The 
cuprous  compounds  contain  a  larger  proportion  of  copper 
than  the  corresponding  cupric  compounds,  e.g.  cuprous 


308  CHEMISTRY 

oxide  is  Cu2O  and  cupric  oxide  CuO.  The  valence  of 
copper  is  I  in  cuprous  compounds  and  II  in  cupric.  Solu- 
ble copper  compounds  are  more  or  less  poisonous.  Cook- 
ing utensils  made  of  copper  should  be  used  with  care. 
Vegetables,  acid  fruits,  and  preserves,  if  boiled  in  them, 
should  be  removed  as  soon  as  cooked.  The  vessels  them- 
selves should  be  kept  bright,  to  prevent  the  formation 
of  copper  salts,  which  might  contaminate  the  contents. 

401.  Copper  Sulphate   or  Cupric   Sulphate,  CuSO4,  is 
a  blue  solid,  and  is  called  blue  vitriol  or  "blue  stone." 
The  crystallized  salt  (CuSO4.5H2O)  effloresces;   heated  to 
240°  C.,  all  the  water  escapes,  leaving  a  whitish  powder 
called  anhydrous  copper  sulphate.     Copper  sulphate  solu- 
tions have  an  acid  reaction  owing  to  hydrolysis  (165). 
Copper  sulphate  is  used  in  electric  batteries,  in  making 
other  copper  salts,  in  calico  printing,  dyeing,  copper  plat- 
ing and  electrotyping.     It  is  poisonous  and  is  one  ingredi- 
ent of  certain  mixtures,  such  as  Bordeaux  mixture,  which 
are  sprayed  upon  trees  to  kill  insects.     Copper  sulphate 
may  be  prepared  by  treating  copper  with  sulphuric  acid 
or  by  oxidizing  copper  sulphide.     Both  methods  are  used. 
Some  of  the  copper  sulphate  of  commerce  is  a  by-product 
obtained  in  refining  gold  and  silver  with  sulphuric  acid. 

402.  Other  Copper  Compounds.  —  Cuprous  oxide  (Cu2O)  is  the 
red  mineral  cuprite.     It  is  precipitated  as  a  reddish  powder  by  heat- 
ing Fehling's  solution  (i.e.  a  mixture  of  solutions  of  copper  sulphate, 
Rochelle  salt,  and  sodium  hydroxide)  with  glucose;    its  formation 
serves  as  a  test  for  glucose  and  sugars  like  it  (223).     Cupric  oxide 
(CuO)  is  a  black  solid  formed  by  heating  copper  in  air.     It  is  used 
to  remove  sulphur  compounds  from    petroleum.     Copper    nitrate, 
Cu(NO3)2,  is  a  blue,  crystallized  solid,  formed  by  the  interaction  of 
copper  and  dilute  nitric  acid.  .  It  is  a  cupric  salt.     It  is  deliquescent, 
very  soluble  in  water,  and  is  readily  decomposed  by  heat  into  cupric 
oxide  (CuO)   and  oxides  of  nitrogen.     Cuprous  sulphide,  Cu2S,  is 


COPPER  —  SILVER  —  GOLD  309 

the  bluish  black  mineral  chalcocite.  Cupric  sulphide,  CuS,  is  the 
black  precipitate  formed  by  passing  hydrogen  sulphide  gas  into  a 
solution  of  a  cupric  salt.  Malachite  is  a  bright  green  mineral  and 
is  often  used  as  an  ornamental  stone.  Azurite  is  a  magnificent  blue, 
crystallized  mineral.  Both  are  basic  carbonates  and  are  ores  of 
copper  (394). 

Silver 

403.  Occurrence  of  Silver.  —  Native  silver  is  some- 
times found  as  flakes  or  wire  in  certain  kinds  of  rock.     The 
chief  ores  are  the  sulphides.     The  simple  sulphide  (silver 
glance,  argentite,  Ag-2S)  is  the  richest  ore  and  is  found  in 
many  localities  in  the  United  States;    complex  ores  con- 
taining antimony  are  widely  distributed.     Small  quanti- 
ties of  native  silver  chloride  (horn  silver,  cerargyrite,  AgCl) 
are  also  found;   it  resembles  wax  or  horn,  and  melts  in  a 
candle  flame.     Alloys  of  silver  with  gold  and  copper  are 
found,  average  California  gold  containing  about  12  per 
cent  silver.     Most  ores  of  lead  and  copper  contain  silver, 
and  this  argentiferous  (or  silver-bearing)  ore  is  one  of  the 
chief  sources  of  silver. 

404.  Metallurgy  of  Silver.  —  Silver  is  extracted  from 
its  ores  by  two  principal  processes.     Ores  containing  free 
silver  or  silver  compounds  that  can  be  easily  changed  into 
free  silver  are  treated  by  the  amalgamation  process.     The 
powdered  ore  is  first  changed  into  silver  chloride  by  roast- 
ing it  with  sodium  chloride.    The  mass  is  then  reduced  to 
silver  by  agitation  with  water  and  iron  (or  an  iron  com- 
pound) ;   the  simplest  equation  for  this  reaction  is  - 

2AgCl      +  Fe  =  2Ag  +      FeCl2 

Silver  Chloride         Iron         Silver         Iron  Chloride 

The  silver  is  removed  by  adding  mercury,  which  forms 
an  amalgam  with  the  silver,  but  not  with  the  other 


3io  CHEMISTRY 

substances.  When  the  amalgam  is  heated,  the  mercury 
distils  off  and  the  silver  —  with  some  gold  —  remains 
behind.  Lead  ores  containing  silver  are  treated  by  the 
Parkes  process.  After  the  sulphur,  arsenic,  and  other 
impurities  have  been  removed  from  the  lead  ores  by  roast- 
ing, the  final  mixture  of  lead,  silver,  and  gold  is  melted 
with  about  i  per  cent  of  zinc.  As  the  mixture  cools,  an 
alloy  of  silver,  gold,  zinc,  and  a  little  lead  rises  to  the 
surface,  solidifies,  and  is  skimmed  off.  This  process  is 
repeated  several  times.  The  skimmings  are  heated  in 
a  retort  to  remove  the  zinc  and  then  in  a  shallow  furnace 
to  remove  the  lead.  The  alloy  of  silver  and  gold  left  in 
the  furnace  can  be  separated  by  two  processes.  In  the 
older  process  the  alloy  is  boiled  with  concentrated  sul- 
phuric acid;  the  gold  is  not  acted  upon,  but  the  silver 
forms  soluble  silver  sulphate.  The  silver  is  displaced 
from  the  solution  by  metallic  copper,  the  latter  forming 
copper  sulphate  (401).  In  the  newer  process  silver  is 
refined  by  electrolysis  in  much  the  same  way  as  copper; 
the  solution  used  is  silver  nitrate  in  nitric  acid  (396). 

405.  Properties  of  Silver.  —  Silver  is  a  lustrous,  white 
metal.  It  is  harder  than  gold-,  but  softer  than  copper. 
Like  copper,  it  is  ductile  and  malleable,  and  may  be  easily 
made  into  various  shapes.  Silver  has  a  specific  gravity  of 
about  10.5,  being  heavier  than  copper,  but  lighter  than  lead. 
It  melts  at  about  960°  C.  Silver  conducts  electricity  the 
best  of  all  the  metals,  but  it  is  too  expensive  for  general 
use.  It  does  not  tarnish  in  air  unless  sulphur  compounds 
are  present,  and  then  the  familiar  black  film  of  silver  sul- 
phide is  produced.  This  blackening  is  especially  noticed 
on  silver  spoons  which  have  been  put  into  eggs  or  mustard, 
and  on  silver  coins  which  have  been  carried  in  the  pocket, 
the  sulphur  in  the  latter  case  coming  from  sulphur  com- 


COPPER  —  SILVER  —  GOLD  3 1 1 

pounds  in  the  perspiration;  the  tarnishing  of  household 
silver  is  due  to  sulphur  compounds  in  illuminating  gas 
or  gas  from  burning  coal.  So-called  " oxidized"  silver  is 
not  oxidized,  but  coated  with  silver  sulphide.  Silver 
is  only  very  slightly  acted  upon  by  hydrochloric  acid, 
and  not  at  all  by  molten  sodium  hydroxide,  potassium 
hydroxide,  or  potassium  nitrate.  Nitric  acid  and  hot 
concentrated  sulphuric  acid  change  it  into  the  nitrate 
(AgNO3)  and  the  sulphate  (Ag2SO4)  respectively,  as  in 
the  case  of  copper.  Potassium  cyanide  in  the  pres- 
ence of  air  and  water  changes  it  into  potassium  silver 
cyanide  (KAg(CN)2). 

Tarnished  silverware  can  be  safely  and  quickly  cleaned  by  an 
electrolytic  process.  A  piece  of  metallic  aluminium  and  the  tarnished 
object  are  immersed  for  a  few  minutes  in  a  hot  solution  of  sodium 
bicarbonate  and  sodium  chloride;  as  soon  as  the  cleaning  is  accom- 
plished, the  object  is  removed,  thoroughly  washed  in  clean,  hot  water, 
and  dried.  The  proportions  for  household  use  are  a  teaspoonful 
each  of  baking  soda  and  common  salt  to  a  quart  of  water.  The  best 
results  are  obtained  when  the  solution  is  very  hot  and  the  two  metals 
are  in  good  contact.  (See  Part  II,  Exp.  192.) 

406.  Alloys  of  Silver.  —  Pure  silver  is  too  soft  for  con- 
stant use,  and  is  usually  hardened  by  adding  a  small 
amount  of  copper.     These  alloys  are  used  as  coins  and 
for  jewelry.     The  silver  coins  of  the  United  States  and 
France   contain   900  parts   of   silver   to    100   of   copper, 
and  are  called  900  fine.     British  silver  coins  are  925  fine; 
this  quality  is  called  sterling  silver,  and  from  it  much 
ornamental  and  useful  silverware  is  made. 

407.  Silver  Plating.  —  Metals  cheaper  than  silver  may  be  coated 
or  plated  with  pure  silver  precisely  as  in  the  case  of  copper.     Plated 
silverware  has  the  appearance  of  solid  or  pure  silver.     The  object  to 
be  plated  is  carefully  cleaned,  and  made  the  cathode  in  a  solution 


312 


CHEMISTRY 


of  potassium  silver  cyanide;   the  anode  is  a  plate  of  pure  silver  (Fig. 

82).  When  the  electric  current  is  passed  the  silver  dissolves  from 

the  anode  and  deposits  upon 
the  cathode.  The  deposit 
of  silver  is  dull,  but  can  be 
brightened  by  rubbing. 


408.    Compounds    of 

'Silver.  --  The  most  im- 
Fig.  82.-  Apparatus  for  Silver  Plating.     pQrtant      compOund     ;s 

silver  nitrate  (AgNOs).  It  is  a  white  crystalline  solid, 
made  by  dissolving  silver  in  nitric  acid,  the  equation  for 
the  reaction  being  - 

3Ag  +   4HN03   =      3AgNOs    +        NO         +  H2O 

Silver        Nitric  Acid        Silver  Nitrate        Nitric  Oxide  Water 

It  is  very  soluble  in  water.  It  turns  dark  in  contact  with 
organic  matter  owing  to  its  reduction  to  metallic  silver. 
For  this  reason  it  blackens  the  skin;  if  applied  long 
enough,  it  disintegrates  the  flesh,  and  is  often  used  by 
physicians  for  this  purpose.  Silver  nitrate  is  the  essen- 
tial substance  used  in  making  indelible  inks;  the  cloth 
to  which  the  ink  is  applied  reduces  the  silver  compound 
to  black  metallic  silver.  Silver  chloride  (AgCl)  is  made 
by  adding  hydrochloric  acid  or  the  solution  of  any  chloride 
to  a  solution  of  a  silver  compound.  Thus  formed,  it  is 
a  white,  curdy  solid,  which  turns  violet  in  the  light  and 
finally  black.  It  is  converted  by  ammonium  hydroxide 
into  a  soluble  complex  compound  (Ag(NH3)2Cl).  The 
formation  and  properties  of  silver  chloride  constitute  the 
test  for  silver.  Silver  bromide  (AgBr)  and  silver  iodide 
(Agl)  are  analogous  to  silver  chloride  in  their  properties 
and  methods  of  formation;  they  are  used  in  photography. 

409.     Photography  is  based  mainly  on  the  fact  that  silver  salts, 
especially  the  bromide  and  iodide,  darken  when  exposed  to  the  light. 


COPPER  —  SILVER  —  GOLD  313 

The  photograph  is  taken  on  a  glass  plate,  coated  on  one  side  with  a 
thin  layer  of  gelatin  containing  very  finely  divided  silver  bromide. 
Sometimes  a  sheet  of  sensitized  celluloid,  called  a  film,  is  used.  The 
plate  or  film  is  exposed  in  the  camera.  The  light  that  comes  from 
the  object  being  photographed  changes  the  silver  salt  in  proportion 
to  its  intensity.  The  exposed  plate,  however,  shows  no  change  until 
it  is  developed.  This  process  consists  in  immersing  the  plate  in  a 
solution  of  a  reducing  agent,  e.g.  hydroquinone,  pyrogallic  acid,  or 
special  mixtures.  As  the  developer  acts  upon  the  silver  salt  on  the 
plate,  the  image  appears.  This  is  really  a  deposit  of  finely  divided 
silver  which  varies  in  thickness  in  proportion  to  the  light  that  fell 
upon  the  plate,  being  thickest  where  the  light  was  most  intense. 
After  the  plate  has  been  properly  developed,  it  still  contains  some 
silver  salt  not  altered  by  the  light;  and  if  the  salt  were  left  on  the 
plate,  the  image  would  be  clouded  and  finally  obliterated  by  the 
light.  The  image  is,  therefore,  fixed  by  dissolving  the  silver  salt  with 
a  solution  of  sodium  thiosulphate  (or  "hyposulphite")  and  washing 
the  plate  thoroughly.  On  the  fixed  plate  the  dark  parts  of  the  object 
appear  light  and  the  light  parts  dark;  and  since  the  image  is  the 
reverse  of  the  object,  the  plate  in  this  condition  is  called  a  negative. 
The  print  is  made  by  laying  sensitized  paper  upon  the  negative  and 
exposing  to  the  sunlight,  so  that  the  light  must  pass  through  the  nega- 
tive first.  Since  the  negative  obstructs  the  light  in  proportion  to  the 
thickness  of  the  silver  deposit,  the  photograph  will  have  approxi- 
mately the  same  shading  as  the  object.  On  some  kinds  of  paper  the 
image  appears  at  once,  but  on  others  it  must  be  developed.  Sub- 
sequent treatment  involves  fixing. 

Gold 

410.  Occurrence.  —  Gold  is  widely  distributed,  but 
not  abundantly  in  many  places.  Unlike  copper  and  silver, 
its  compounds  are  few  and  rare;  the  only  important  ones 
are  the  tellurides  (compounds  of  tellurium,  e.g.  (AuAgTe2) 
found  in  Colorado.  It  is  never  found  pure,  being  alloyed 
with  silver  and  occasionally  with  copper  or  iron  (403) .  It 
is  disseminated  in  fine,  almost  invisible,  particles  among 
ores  of  other  metals,  though  not  so  abundantly  as  silver. 


314  CHEMISTRY 

Much  gold  is  found  in  veins  of  quartz,  and  in  the  sand 
formed  by  the  disintegration  of  gold-bearing  rocks. 

The  chief  gold-producing  countries  are  the  United  States,  Aus- 
tralia, South  Africa,  and  Russia.  The  United  States  produces  annu- 
ally over  four  million  troy  ounces,  which  come  largely  from  Colorado, 
California,  and  Alaska. 

411.  Mining,  Metallurgy,  and  Refining.  —  Gold  is  sometimes 
mined  by  washing  the  gold-bearing  sand  in  large  pans  or  cradles. 
In  placer  mining  and  hydraulic  mining  streams  of  water  wash  away 
the  lighter  materials,  but  leave  the  heavier  gold  behind.  From  this 
mixture,  gold  and  silver  are  extracted  with  mercury.  In  vein  or 
quartz  mining  the  lumps  of  gold-bearing  quartz  are  crushed  to  a  fine 
powder  in  stamp  mills;  the  moistened  mass  is  passed  over  copper 
plates  coated  with  mercury  which  collects  or  dissolves  the  gold  and 
other  metals.  The  amalgam  is  heated,  as  in  the  metallurgy  of 
silver,  to  remove  the  mercury,  and  the  gold  is  extracted  from  the 
residue  (404).  In  the  chlorination  process  the  crushed  ore  is  roasted 
and  then  treated  with  chlorine  or  with  bleaching  powder  and  sul- 
phuric acid;  this  operation  forms  a  soluble  gold  chloride  (AuCl3), 
from  which  the  gold  is  precipitated  as  a  fine  powder  by  ferrous  sul- 
phate or  other  reducing  agents.  In  the  cyanide  process  the  crushed 
ore,  or  the  slime  from  a  previous  extraction,  is  mixed  with  a  weak 
solution  of  potassium  cyanide  and  exposed  to  the  air;  this  operation 
changes  the  gold  into  a  soluble  cyanide,  the  equation  for  the  reaction 
being  — 

4Au  +  8KCN    +    02    +  2H20  =  4KAu(CN)2  +  4KOH 

Gold         Potassium         Oxygen        Water  Gold  Potassium         Potassium 

Cyanide  Cyanide  Hydroxide 

The  gold  is  separated  from  this  solution  by  electrolysis  or  by  treat- 
ment with  zinc. 

Refining  is  accomplished  by  parting  or  by  electrolysis.  By  the 
old  parting  process  known  as  quartation  an  alloy  of  gold  and  silver, 
in  which  the  gold  is  about  one  fourth  of  the  whole,  is  treated  with 
nitric  acid;  this  operation  changes  the  silver  into  the  nitrate  from 
which  the  pure  gold  may  be  readily  removed.  The  metals  may  be 
parted  by  the  method  described  under  silver,  viz.  by  boiling  with 
concentrated  sulphuric  acid.  By  this  treatment  the  silver  is  dis- 
solved and  the  gold,  which  is  about  one  sixth  of  the  alloy,  is  left  as  a 


COPPER  —  SILVER  —  GOLD  315 

brownish,  porous  mass.  In  an  electrolytic  process  of  separation, 
which  is  now  used  in  the  United  States  mints,  the  electrolytic  solu- 
tion is  a  mixture  of  hydrochloric  acid  and  gold  chloride,  the  anode 
is  an  alloy  rich  in  gold,  and  the  cathode  is  pure  gold.  Gold  is  de- 
posited on  the  cathode,  and  the  silver  forms  silver  chloride  around 
the  anode. 

The  purity  of  gold  is  expressed  in  carats.  Pure  gold  is  24  carats 
fine;  an  alloy  containing  22  parts  of  gold  and  2  parts  copper  is  22 
carat  gold.  (Compare  171  footnote.) 

412.  Properties  and  Uses.  —  Gold  is  a  yellow  metal. 
It  is  nearly  as  soft  as  lead,  and  is  the  most  ductile  and 
malleable  of  all  the  metals.     The  leaf  into  which  it  can  be 
beaten  is  very  thin.     Gold  is  one  of  the  heaviest  metals, 
its  specific  gravity  being  about  19.3.     The  melting  point 
is  1063°  C.     It  forms  alloys  with  most  metals.     Air,  oxy- 
gen, and  most  acids  do  not  attack  it,  and  for  this  reason  it 
is   sometimes   called  a  noble  metal.     It  is  changed  into 
gold   chloride  (AuCls)  by  chlorine   and  aqua  regia  (87). 
With    potassium   cyanide,   as   described  above   (411),  it 
forms  potassium  gold  cyanide  (KAu(CN)j). 

Pure  gold  is  too  soft  for  use  as  jewelry  or  coins,  and  it  is 
usually  alloyed  with  copper  or  silver.  Gold  coins  contain 
gold  and  copper.  The  United  States  standard  gold  coins 
contain  90  per  -cent  gold  and  10  per  cent  copper.  Gold 
leaf  of  various  grades  is  used  to  ornament  books  and 
signs.  Jewelers  use  gold  for  many  purposes;  such  gold 
varies  from  12  to  22  carats  in  purity.  On  account  of  its 
malleability,  feeble  chemical  action,  and  beauty,  gold  is 
used  by  dentists  for  filling  teeth. 

413.  Compounds  of  Gold  are  readily  decomposed  by  metals,  weak 
reducing  agents   (e.g.   ferrous  sulphate  or  hydrogen  sulphide),  fine 
solids  like  charcoal,  and  by  electrolysis.     When  gold  is  dissolved  in 
aqua  regia  and  the  acid  removed  by  evaporation,  the  resulting  gold 
chloride  (AuQ3)  gives  with  stannous  chloride  solution  a  beautiful 


3i6  CHEMISTRY 

purple  precipitate;  the  latter  is  called  "purple  of  Cassius,"  and  is 
finely  divided  gold.  Its  formation  is  the  test  for  gold.  The  process 
of  gold  plating  is  the  same  as  silver  plating;  the  solution  contains  a 
potassium  gold  cyanide  (KAu(CN)2  or  KAu(CN)4)  and  the  anode  is 
gold.  Much  cheap  jewelry  is  gold  plated. 

414.  Displacement  of  Metals. --The  deposition  of 
metallic  copper  and  the  displacement  of  mercury  by 
copper  itself  (397)  are  examples  of  a  kind  of  chemical 
change  in  which  most  metals  can  participate.  The 
displacing  relations  of  metals,  so  to  speak,  have  been 
carefully  studied  and  the  familiar  metals  can  be  arranged 
in  a  series  called  the  - 

ELECTROCHEMICAL  SERIES  OF  THE  METALS 


Magnesium 

Cobalt 

Copper 

Aluminium 

Nickel 

Mercury 

Zinc 

Tin 

Silver 

Cadmium 

Lead 

Platinum 

Iron 

Hydrogen 

Gold 

In  this  series  each  free  metal  displaces  succeeding  metals 
from  their  solutions,  and  is  in  turn  displaced  from  its 
solutions  by  preceding  metals.  For  example,  zinc  (except 
when  very  pure),  being  near  the  beginning,  displaces  most 
of  the  metals  in  the  series,  while  copper,  being  near  the 
end,  displaces  only  a  few. 

Hydrogen  is  not  a  metal  in  the  common  acceptance  of  this  term, 
but  it  is  usually  included  in  the  electrochemical  series  of  the  metals, 
because  all  the  metals  that  precede  hydrogen  displace  it  from  most 
acids,  while  those  that  follow  do  so  rarely,  if  at  all.  That  is,  hydro- 
gen behaves  in  this  respect  like  other  members  of  the  series. 


CO  PPER  —  SILVER  —  GOLD  3 1 7 


EXERCISES 

1.  Describe  the  refining  of  copper  by  electrolysis. 

2.  State  the  physical  properties  of  copper  that  fit  it  for  electrical  uses. 

3.  Describe  several  tests  for  copper. 

4.  Name  eight  alloys  of  copper.     State  the  uses  of  three. 

5.  Describe  the  extraction  of  silver  by  (a)  the  amalgamation  process 
and  (b)  the  Parkes  process. 

6.  Describe  the  process  of  (a)  silver  plating  and  (b)  gold  plating. 

7.  Describe  briefly  the  essential  operations  in  photography. 

8.  What  is  (a)  blue  vitriol,  (b)  argentiferous  lead,  (c)  oxidized  silver, 
(d)  sterling  silver,  («)  coin  silver,  (/)  lunar  caustic,  (g)  "hypo,"  (/?)  fool's 
gold,  (i)  gold  leaf? 

9.  How  is  gold  purified?     What  is  18  carat  gold? 

10.  Practical  topics:  (a)  If  silver  cost  only  ten  cents  a  pound,  for  what 
could  it  be  used?  (b)  Why  is  gold  "valuable"?  (c)  A  "tin"  can  will  dis- 
place copper  from  copper  solutions.  Why?  (d)  Why  are  the  sparks  from 
a  trolley  wire  colored  green?  (e)  Nitric  acid  produces  a  blue  stain  on  a 
silver  coin.  Why?  (/)  A  2  carat  diamond  is  set  in  an  18  carat  gold  ring. 
Explain. 

PROBLEMS 

1.  Calculate  the  weight  of  copper  in  (a)  10  gm.  of  crystallized  copper 
sulphate,  (b)  i  metric  ton  of  malachite,  and  (c)  2000  Ib.  of  chalcocite. 

2.  Calculate  the  weight  of  (a)  silver  in  i  kg.  of  silver  nitrate  and  (b)  of 
gold  in  i  gm.  of  potassium  auricyanide  (KAu(CN)4). 

3.  Calculate  the  percentage  composition  of  the  two  copper  oxides  and 
show  that  they  illustrate  the  law  of  multiple  proportions.     (Use  exact 
atomic  weights.) 

4.  Write  the  formulas  of  the  following  compounds  by  applying  the  prin- 
ciple of   valence:     Silver  sulphide,  silver   hydroxide,  silver  oxide,  cupric 
bromide,  cuprous  chloride,  cupric  cyanide,  auric  oxide,  aurous  bromide, 
auric  chloride,  silver  acetate,  cuprous  oxide,  cupric  sulphate,  cuprous  iodide, 
silver  nitrate,  aurous  oxide,  auric  sulphide,  silver  orthophosphate.     Calcu- 
late the  per  cent  of  the  metal  in  five  of  these  compounds. 

5.  Calculate    the    simplest   formulas    corresponding    to    the    following 
analyses:  (a)  Cu  =  96.72,  H  =  3.82  ;  (b}  Cu  =  57.46,  H  =  0.91,  O  =  36.2, 
C  =  5-435    (c)  Ag  =  87.09,  S  =  12.9;    (d)  Au  =  64.81,  Cl  =  35.09. 

6.  How  many  grams  of  gold  in  a  14  carat  ring  which  weighs  14  gm.? 

7.  How  much  (a)  silver  and  (b)  copper  can  be  obtained  from  an  Ameri- 
can ten  cent  coin  which  weighs  2.44  gm.? 

8.  Complete  and  balance  the  equation  CuSC>4  +  HaS  =  CuS  H . 

How  much  CuS  can  be  obtained  from  24  gm.  of  CuSC>4? 


CHAPTER  XXVI 


CALCIUM  —  STRONTIUM  AND  BARIUM 

These  elements  form  a  natural  family  called  the  alka- 
line earth  metals. 

Calcium 

415.  Occurrence  of  Calcium.  —  Calcium  is  never  found 
free.  Combined  calcium  makes  up  about  3.5  per  cent  of 
the  earth's  crust.  The  most  abundant  compound  is 
calcium  carbonate  (CaCO3).  Many  rocks  are  complex 
silicates  of  calcium  and  other  metals.  The  extensive 
deposits  of  calcium  phosphate  and 
calcium  borate  have  been  mentioned. 
Calcium  sulphate  (CaSO4)  also  oc- 
curs abundantly.  Calcium  com- 
pounds are  essential  to  the  life  of 
plants  and  animals,  being  found  in 
the  leaves  of  plants,  and  in  the 
bones,  teeth,  and  shells  of  animals.. 
Many  rivers  and  springs  contain  cal- 
cium salts,  especially  the  acid  car- 
bonate and  sulphate  (417,  423). 

416.  Preparation  and  Properties.  —  Me- 
tallic calcium  is  prepared  by  the  electroly- 
sis of  melted  calcium  chloride.  One  form 
of  apparatus  is  shown  in  Fig.  83.  The  anode 
is  a  graphite  crucible  (A)  and  the  cathode  is 


Fig.  83.  —  Apparatus  for 
Preparing  Calcium  by 
Electrolysis. 


a  rod  of  iron  (B),  which  dips  into  the  melted  calcium  chloride  a  short 
distance  and  is  so  adjusted  that  it  can  be  elevated  by  a  screw  (C). 
The  lower  part  of  the  crucible,  which  is  kept  cool  by  running  water 


CALCIUM  —  STRONTIUM   AND   BARIUM       319 

(E,  E),  contains  solid  calcium  chloride.  When  the  current  passes, 
calcium  is  deposited  on  the  cathode,  which  is  slowly  raised  so  that 
its  end  is  kept  in  contact  with  the  surface  of  the  melted  chloride;  the 
irregular  rod  of  deposited  calcium  (D)  thus  becomes  the  end  of  the 
cathode. 

Calcium  is  a  silvery  white  metal.  Its  specific  gravity  is  1.55  and 
its  melting  point  is  810°  C.  It  tarnishes  slowly  in  air.  When  heated, 
it  combines  with  most  non-metals.  If  burned  in  air,  it  forms  both 
the  oxide  (CaO)  and  the  nitride  (Ca3N2).  It  interacts  with  water, 
slowly  at  ordinary  temperatures,  rapidly  at  high  temperatures.  The 
equation  for  the  reaction  is  — 

Ca          +       2H20     =        -Ca(OH)2         +        H2 

Calcium  Water  Calcium  Hydroxide  Hydrogen 

It  also  interacts  readily  with  acids,  thus:  — 

Ca     +          2HC1  CaCl2         +       H2 

Calcium  Hydrochloric  Acid         Calcium  Chloride  Hydrogen 

417.  Calcium  Carbonate,  CaCO3.  —  Large  quantities 
are  found  in  many  regions.  The  commonest  form  is 
limestone.  Pure  limestone  is  white  or  gray,  but  impuri- 
ties, especially  organic  matter  and  iron  compounds,  pro- 
duce many  colored  varieties.  Much  limestone  contains 
silica,  clay,  iron  and  aluminium  compounds,  and  the 
fossil  remains  of  plants  and  animals.  Hard,  crystalline 
limestone  which  takes  a  good  polish  is  called  marble; 
it  is  extensively  used  as  a  building  and  an  ornamental 
stone.  Calcite  is  an  abundant  form  of 
crystallized  calcium  carbonate;  a  trans- 
parent variety  called  Iceland  spar  has 
the  property  of  double  refraction,  i.e. 
of  making  objects  appear  double  (Fig.  p.g  84  _ Crystallized 
84).  Different  varieties  of  calcium  car-  Iceland  Spar  Show- 
bonate  are  used  in  making  lime,  cement,  ing  Double  Refrac- 
iron,  glass,  and  sodium  carbonate. 

Calcium  carbonate  is  soluble  in  water  containing  carbon  dioxide, 
owing  to  its  transformation  into  the  soluble  acid  calcium  carbonate 


320  CHEMISTRY 

(H2Ca(C03)2)  (188).  As  this  water  works  its  way  underground  in 
limestone  regions  the  limestone  is  dissolved  and  caves  are  often 
formed  or  enlarged.  When  the  water  enters  a  cave  and  drips  from 
the  top,  the  acid  calcium  carbonate  decomposes  and  calcium  car- 
bonate is  redeposited,  often  forming  stalactites  and  stalagmites. 
The  stalactites  hang  from  the  roof  like  icicles,  and  are  often  exquisitely 
shaped;  the  stalagmites,  which  grow  up  from  the  floor,  sometimes 
meet  the  stalactites  and  form  a  column.  Mexican  onyx  is  a  variety 
of  stalagmite.  Vast  deposits  of  this  beautiful  mineral  are  found  in 
Algeria  and  Mexico.  It  is  translucent  and  delicately  colored,  and  is 
used  as  an  ornamental  stone,  especially  for  altars,  table  tops,  mantels, 
and  lamp  standards.  Travertine  occurs  near  many  springs  in  Italy. 
When  fresh,  it  is  soft  and  porous,  but  it  soon  hardens  and  becomes  a 
durable  building  stone  in  dry  climates.  A  portion  of  the  walls  of  the 
Colosseum  and  St.  Peter's  is  travertine. 

Limestone  often  contains  shells  and  fossils,  confirming  the  belief 
that  limestone  is  the  remains  largely  of  the  shells  of  animals.  The 
calcium  carbonate  dissolved  in  the  ocean  is  transformed  by  marine 
organisms  into  shells  and  bony  skeletons.  The  hard  parts  of  these 
animals  accumulate  in  vast  quantities  on  the  ocean  bottom,  become 
compact,  often  hardened  and  crystallized,  and  subsequently  form  a 
part  of  the  land.  On  the  coast  of  Florida,  coquina  or  shell  rock  is 
found.  It  is  a  mass  of  fragments  of  shells  cemented  by  calcium  car- 
bonate, and  in  time  will  become  compact  limestone.  Chalk  is  the 
remains  of  shells  of  minute  animals.  When  examined  under  a  micro- 
scope, a  good  specimen  is  seen  to  consist  almost  entirely  of  tiny  shells. 

418.  Calcium  Oxide,  CaO,  is  the  familiar  substance 
lime.  It  is  a  hard,  white  solid.  Pure  lime  is  almost 
infusible,  and  when  heated  in  the  oxyhydrogen  flame,  it 
gives  an  intensely  bright  light,  sometimes  called  the  lime 
light  (26).  In  the  intense  heat  of  the  electric  furnace  it 
melts  and  boils.  Lime  containing  impurities,  like  sand, 
clay,  and  iron  compounds,  melts  quite  readily  into  a  glass 
or  slag.  Exposed  to  the  air,  lime  becomes  "air  slaked," 
i.e.  it  slowly  absorbs  water  and  carbon  dioxide,  swells, 
and  soon  crumbles  to  a  powder,  which  is  a  mixture  of  cal- 


CALCIUM  —  STRONTIUM   AND   BARIUM       321 


cium  hydroxide  and  calcium  carbonate.  Lime  and  water 
combine  readily;  considerable  heat  is  liberated,  as  is  often 
seen  when  mortar  is  being  prepared.  This  operation  is 
called  "slaking,"  and  the  product  is  "slaked  lime."  The 
equation  for  the  reaction  is  - 

CaO       +  H2O  =       Ca(OH)2       +15,540  calories 

Calcium  Oxide        Water        Calcium  Hydroxide 

Sometimes  water  leaks  into  cars  or  buildings  in  which 
lime  is  stored  and  the  heat  evolved  causes  a  serious  fire. 

Lime  is  used  in  preparing  mortar,  cement,  and  metals, 
in  making  bleaching  powder,  calcium  carbide,  sodium  hy- 
droxide, and  glass,  in  purifying  illuminating  gas  and  sugar, 
in  removing  hair  from  hides,  in  dyeing  and  bleaching 
cotton  cloth,  and  as  a  disinfectant  and  fertilizer. 

Lime  is  manufactured  by  heating  limestone  in  a  partly  closed  cavity 
or  vessel.  The  decomposition  takes  place  according  to  the  equation — 

CaC03  CaO         +        CO2 

Calcium  Carbonate        Calcium  Oxide        Carbon  Dioxide 

The  carbon  dioxide  gas  escapes  and 
the  lime  is  left  in  the  kiln. 

Limestone  was  formerly 
"burned"  in  a  cavity  on  a  hillside, 
and  in  some  regions  it  is  so  pre- 
pared to-day.  An  arch  of  lime- 
stone is  built  across  the  cavity 
above  the  fire  pit,  and  limestone  is 
introduced  until  the  kiln  is  full. 
These  kilns  are  being  replaced  by 
modern  kilns  (Fig.  85),  which  are 
constructed  so  that  the  heat  can 
be  regulated  (at  B,  B),  the  gases  @ 
swept  out,  the  raw  material  intro- 
duced (at  A),  and  the  product  re- 
moved continuously  (at  C,  C).  Fig.  85.  —  Continuous  Limekiln. 


322  CHEMISTRY 

419.  Cement.  —  Pure  or  nearly  pure  limestone  yields 
a  product  called  quicklime  (or  simply  lime)  because  it  acts 
quickly  with  water  in  contrast  with  lime  which  is  impure 
or  partly   slaked.     Limestone   containing   about   10  per 
cent  of  clay  and  silica  yields  a  mixture  of  calcium  silicate 
and  aluminate  besides  lime.     Unlike  quicklime  it  slakes 
very  slowly  and  hardens  under  water  as  well  as  in  air.     It 
is  therefore  called  hydraulic  lime.     Hydraulic  lime  is  in- 
termediate between  lime  and  cement.     Cement  is  made 
from   natural   or   artificial   mixtures   of   limestone,   clay, 
sand,  and  iron  oxide.     Portland  cement  is  manufactured 
by  heating  the  pulverized,  carefully  proportioned  mixture 
in  long  rotatory  kilns  (Fig.  86) .     The  mixture  enters  at  C 

and    gradually 
t    works    its    way 
along  through  the 
slowly    rotating 

Fig.  86. -Cement  Kiln.  kiln     against     the 

flames    and    hot 

gases  produced  by  burning  coal  dust,  which  is  forced  into 
the  kiln  at  A  by  a  powerful  air  blast.  The  mixture  inter- 
acts and  forms  a  semifused,  gray-black  mass  called  clinker, 
which  drops  out  at  B,  and  is  subsequently  ground  to  a  very 
fine  powder.  Ground  gypsum  is  often  added.  A  mixture 
of  cement,  sand,  water,  and  crushed  stone  is  known  as 
concrete,  which,  as  well  as  cement  itself,  is  now  exten- 
sively used  as  a  construction  material.  Sometimes  con- 
crete is  strengthened  by  imbedding  rods  of  iron  or  steel 
in  it,  and  it  is  then  called  re-enforced  concrete.  (See 
Part  II,  Exp.  197.) 

420.  Calcium  Hydroxide,  Ca(OH)2,  is  a  white  powder. 
It  is  sparingly  soluble  in  water,  but  more  soluble  in  cold 
than  in  warm  water.     The  solution  has  a  bitter  taste  and 


CALCIUM  — STRONTIUM   AND   BARIUM       323 

an  alkaline  reaction;  it  is  called  limewater  and  is  often 
used  as  a  medicine.  Exposed  to  the  air,  limewater  be- 
comes covered  with  a  thin  crust  of  calcium  carbonate, 
owing  to  the  absorption  of  carbon  dioxide.  For  the 
same  reason,  limewater  becomes  milky  or  cloudy  when 
carbon  dioxide  is  passed  into  it.  The  formation  of  calcium 
carbonate  in  this  way  is  the  test  for  carbon  dioxide  (183). 

Limewater  is  prepared  by  carefully  adding  lime  to  considerable 
water,  allowing  the  mixture  to  stand  in  a  stoppered  bottle  until  the 
solid  has  settled,  and  then  removing  the  pure  liquid.  When  con- 
siderable calcium  hydroxide  is  suspended  in  the  liquid,  the  mixture 
is  called  milk  of  lime.  Ordinary  whitewash  is  thin  milk  of  lime. 

421.  Mortar  is  a  thick  paste  formed  by  mixing  lime,  sand,  and 
water.     It  slowly  hardens  or  "sets,"  owing  to  the  loss  of  water  and 
to  the  absorption  of  carbon  dioxide.     It   hardens  without  much 
shrinking,  and  when  placed  between  bricks  or  stones  holds   them 
firmly  in  place.     The  sand  makes  the  mass  porous  and  thus  facil- 
itates  the  change  of  the  hydroxide  into  the  carbonate.     Hair  is 
sometimes  added  to  make  the  mortar  stick  better,  especially  when 
it  is  used  as  plaster  for  walls. 

422.  Calcium  Sulphate,  CaSCV  —  Extensive    deposits 
of  the  different  forms  of  calcium  sulphate  are  found  in 
many  localities;  in  the  United  States  large  quantities  are 
obtained  in  New  York,  Michigan,  and  the  middle  West. 
Gypsum  is  the  commonest  form;  it  occurs  as  white  masses 
which  have  the  composition  CaSO^H^O.     A  translucent 
variety  of  gypsum  is  called  selenite.    The  mineral  anhydrite 
is  anhydrous  calcium  sulphate  (CaSC^).     Gypsum  is  used 
as  an  ingredient  of  some  fertilizers  and  in  making  plaster 
of  Paris,  paper,  white  paint,  and  cement. 

Plaster  of  Paris  is  a  fine  white  powder  made  by  heating  gypsum 
to  the  proper  temperature  (about  125°  C.).  This  powder  is  essen- 
tially a  compound  having  the  composition  (CaSO4)2.H2O.  If  mois- 
tened with  water,  it  swells  and  quickly  sets  or  solidifies  to  a  hard  mass 


324  CHEMISTRY 

which  consists  of  a  network  of  very  small  crystals.  The  equation 
for  the  setting  of  plaster  of  Paris  may  be  written  — 

(CaS04)2.H2O     +     3H20     =     2(CaSO4.2H2O) 

Pkster  of  Paris  Water  Gypsum 

Plaster  of  Paris  is  used  to  coat  plastered  walls,  to  cement  glass  to  metal, 
but  more  largely  to  make  casts  and  reproductions  of  statues  and  small 
objects.  Stucco  is  essentially  a  mixture  of  glue  and  plaster  of  Paris. 

423.    Calcium  Compounds  and  Hardness  of  Water.  - 

Calcium  sulphate  is  slightly  soluble  in  water,  and  calcium 
carbonate,  as  we  have  already  seen,  is  changed  into  a 
soluble  acid  carbonate  by  water  containing  carbon  dioxide. 
Water  containing  these  and  other  dissolved  salts  of  cal- 
cium is  called  hard  water.  They  form  sticky,  insoluble 
compounds  with  soap,  and  as  long  as  water  contains  such 
salts,  the  soap  is  useless  as  a  cleansing  agent.  More- 
over hard  water  if  used  in  boilers  forms  deposits  on  the 
inside  of  the  boiler,  thereby  causing  waste  of  heat;  in  some 
cases  acids  are  liberated  which  corrode  the  boiler.  Heat 
decomposes  acid  calcium  carbonate,  and  the  hardness  due 
to  acid  calcium  carbonate  is  called  temporary  hardness, 
because  boiling  removes  it.  Temporary  hardness  can 
also  be  removed  by  adding  enough  calcium  hydroxide  to 
change  the  acid  calcium  carbonate  into  normal  calcium 
carbonate.  Boiling  does  not  affect  calcium  sulphate,  and 
water  containing  this  salt  is  said  to  have  permanent  hard- 
ness. Magnesium  sulphate  and  chloride,  like  the  corre- 
sponding calcium  salts,  produce  permanent  hardness. 
Permanently  hard  water  can  be  softened,  however,  by 
adding  sodium  carbonate,  which  converts  the  calcium 
and  magnesium  salts  into  insoluble  carbonates,  thus:  - 

CaSO4      +      Na2CO3     =      CaCO3        +     Na2SO4 

Calcium  Sulphate    Sodium  Carbonate   Calcium  Carbonate   Sodium  Sulphate 

(See  Part  II,  Exp.  199.) 


CALCIUM  —  STRONTIUM  AND   BARIUM       325 

424.  Calcium  Chloride,  CaCl2,  is  a  white  solid.     It  absorbs  mois- 
ture, and  is  used  to  dry  gases  and  liquids  (61).     Calcium  chloride  is 
found  in  small  quantities  in  some  of  the  Stassfurt  salts  (377).     It  is 
a  by-product  in  the  manufacture  of  sodium  carbonate  by  the  Solvay 
process  (366).     A  solution  of  calcium  chloride  is  used  as  a  brine  in  the 
manufacture  of  ice  (104). 

425.  Other    Calcium    Compounds.  —  Important    calcium    com- 
pounds already  described  are  the  fluoride,  carbide,  phosphates,  and 
hypochlorite.     Calcium  sulphide  (CaS)  is  formed  by  reducing  cal- 
cium sulphate  with  carbon;    like  other  sulphides,  it  stains  silver 
brown.     Calcium  oxalate  (CaC2O4)  is  a  white  solid  formed  by  the  inter- 
action of  ammonium  oxalate  and  a  dissolved  calcium  compound;   it 
is  insoluble  in  acetic  acid  but  soluble  in  hydrochloric  acid.     Its  for- 
mation and  properties  serve  as  a  test  for  calcium.     Another  test  for 
calcium  is  the  light  red  color  imparted  to  the  Bunsen  flame.     Calcium 
nitrate  (Ca(NO3)2)  and  calcium  cyanamide  (CaN2C)  are  made  from 
the  nitrogen  of  the  atmosphere  and  are  used  as  fertilizers  because 
they  provide  nitrogen  in  a  form  easily  taken  up  by  plants.     The  com- 
mercial substances  are  dark  solids.     For  the  nitrate  see  106.     The 
cyanamide  is  made  by  passing  nitrogen  over  very  hot  calcium  car- 
bide, the  equation  for  the  reaction  being  — 

N2         +        CaC2  CaN2C         +        C 

Nitrogen  Calcium    Carbide      Calcium    Cyanamide  Carbon 

Strontium  and  Barium 

426.  Strontium,   Sr,   and  Barium,  Ba,   are  uncommon  metallic 
elements.     The  metals  themselves  never  occur  free,  and  are  hardly 
more  than  chemical  curiosities.     Their  compounds  resemble  those  of 
calcium;  some  are  useful. 

427.  Compounds   of   Strontium.  —  The  important   native   com- 
pounds are  the  beautifully  crystallized  minerals,  strontianite  (stron- 
tium carbonate,  SrCO3)  and  celestite  (strontium  sulphate,  SrSO4). 
Strontium  oxide  (strontia,  SrO),  like  lime,  is  made  by  heating  the 
carbonate.     It    unites   with    water    to    form    strontium    hydroxide 
(Sr(OH)2),  which  is  used  in  refining  beet  sugar.     Strontium  nitrate 
(Sr(NO3)2)  and  other  salts  of  strontium  color  a  flame  crimson;    the 
nitrate  is  used  in  making  red  signal  lights  and  fireworks,  especially 
red  fire.     The  latter  is  essentially  a  mixture  of  potassium  chlorate, 
shellac,  and  strontium  nitrate.     The  production  of  the  crimson  colored 


326  CHEMISTRY 

flame  is  a  test  for  strontium.  Another  test  is  the  precipitation  of 
white  strontium  sulphate  by  the  addition  of  calcium  sulphate  solution 
to  the'solution  of  a  strontium  salt.  (See  Part  II,  Exp.  200.) 

428.  Compounds  of  Barium.  —  The  most  abundant  native 
compounds  are  witherite  (barium  carbonate,  BaCO3)  and  barite 
(barium  sulphate,  BaSO4).  Barium  hydroxide  (Ba(OH)2)  solution  is 
often  called  baryta  water,  and  like  limewater  it  forms  an  insoluble  car- 
bonate (BaCO3)  when  exposed  to  carbon  dioxide.  Barium  chloride 
(BaCl2)  is  used  to  test  for  sulphuric  acid  and  soluble  sulphates,  be- 
cause it  readily  interacts  with  them  and  forms  insoluble  barium 
sulphate  (BaSO4);  conversely,  this  serves  as  the  test  for  barium. 
This  precipitated  salt  is  a  fine,  white  powder,  and  being  cheap  and 
heavy  it  is  a  common  adulterant  of  ordinary  white  paint.  Ground 
native  barium  sulphate,  often  called  barytes,  has  a  similar  use. 
Barium  sulphate  is  also  used  to  increase  the  weight  of  paper  and  to 
give  it  a  gloss.  Barium  salts  color  a  flame  green,  and  barium  nitrate 
(Ba(NO3)2)  is  used  in  making  fireworks,  especially  green  fire.  The 
production  of  the  green  flame  is  a  test  for  barium.  Commercial 
barium  sulphide  (BaS),  as  well  as  the  sulphides  of  calcium  and  stron- 
tium, shine  feebly  in  the  dark,  after  having  been  exposed  to  a  bright 
light.  On  account  of  this  property  they  are  used  in  making  luminous 
paint.  Barium  chromate  (BaCrO4)  is  obtained  as  a  yellow  precipi- 
tate by  the  interaction  of  potassium  dichromate  (or  potassium  chro- 
mate) and  a  soluble  barium  compound;  being  a  colored  compound, 
its  formation  is  sometimes  used  as  a  test  for  barium.  (See  Part  II, 
Exp.  201.) 

EXERCISES 

1.  Name  several  native  compounds  of  calcium.     What  proportion  of 
the  earth's  crust  is  calcium?     Compare  this  proportion  with  that  of  other 
abundant  elements. 

2.  Review  topics:    (a)  The  properties  of  normal  and  acid  calcium  car- 
bonate,    (b)  Calcium    compounds    previously    studied,     (c)  Compare    the 
manufacture  of  calcium  and  sodium. 

3.  Practical    topics:     (a)  Does   limestone   "burn"?     (&)  Is   the   term 
limewater  accurate?    Why?     (c)  How  should  lime  be  stored?     (d)  Compare 
the  setting  of  plaster  of  Paris  and  mortar,     (e)  Suggest  experiments  to 
show  that  lime  is  calcium  oxide. 

4.  State  the  properties  and  uses  of  lime.     How  is  it  made?     Discuss 
the  manufacture  of  cement. 


CALCIUM  —  STRONTIUM   AND   BARIUM       327 

5.  Starting  with  calcium,  how  would  you  prepare  successively  the 
oxide,  hydroxide,  carbonate,  chloride,  and  metal? 

6.  What  is  plaster  of  Paris?     Why  so  called?     How  is  it  prepared? 
What  is  its  chief  property?     What  are  its  uses?     What  is  the  chemical 
explanation  of  "setting"?     What  is  stucco? 

7.  What  is  hard  water?     How  does  it  act  with  soap?     What  is  (a)  tem- 
porary hardness  and  (b)  permanent  hardness?     How  can  each  be  removed? 
What  is  soft  water?     Why  is  rain  water  often  called  soft  water? 

8.  Essay  topics:    (a)   Famous  limestone  caves,     (b)   The  cement  in- 
dustry,    (c)  Colored  fireworks,     (d)  Industrial  uses  of  lime.      (e)  Chalk. 
(/)  Coral,     (g)  Uses  of  barium  and  strontium  compounds. 

9.  Write  equations  for  the  reactions  necessary  to  prepare  (a)  barium 
nitrate  from  barium  carbonate,  (b)  barium  hydroxide  from  barium  chloride, 
(c)  strontium  carbonate  from  strontium  hydroxide. 

10.   State  the  tests  for  (a)  calcium,  (b)  strontium,  and  (c)  barium. 

PROBLEMS 

1.  How  many  grams  of  calcium,  strontium,  or  barium  can  be  obtained 
from  (a)  150  gm.  of  calcium  chloride,  (b)  i  metric  ton  of  Iceland  spar,  (c) 
250  gm.  of  SrSO4,  (d)  27  gm.  of  BaCO3? 

2.  Calculate  the  percentage  composition  of  the  two  oxides  of  barium 
and  show  that  they  illustrate  the  law  of  multiple  proportions.     (Use  exact 
atomic  weights.) 

3.  Calculate  the  simplest  formulas  from  the  following  data:    (a)  Ca  = 
29.49,  O  =  46.92,  S  =  23.59;    (b)  6.87  gm.  of  Ba  unite  with  1.6  gm.  of  O; 
(c)  Sr  =  72.01,  O  =  26.34,  H  =  164. 

4.  Write   the  formulas   of  the  following  compounds  by  applying  the 
principle   of   valence:    Calcium   chlorate,   calcium  permanganate,  calcium 
fluoride,  calcium  silicate,  calcium  sulphide,  barium  iodide,  barium  chromate, 
barium  silicate,  strontium  monoxide,  strontium  chlorate,  strontium  dichro- 
mate.    Calculate  the  per  cent  of  the  metal  in  any  three  of  these  compounds. 

5.  Calculate  the  atomic  weight  of  calcium,  strontium,  or  barium  from: 
(a)  31.20762  gm.  of  CaCO3  give  17.49526  gm.  of  CaO;   (b)  0.5  gm.  of  SrCO3 
give  75.54  cc.  of  CO2  (standard  conditions);   (c)  100  gm.  of  BaCl2  give  112.1 
gm.  of  BaSO-j. 

6.  Express  the  following  reactions  by  equations:    (a)  Carbon  dioxide, 
water,  and  calcium  carbonate  form  acid  calcium  carbonate;    (b)  barium 
hydroxide  and  carbon  dioxide  form  barium  carbonate  and  water;   (c)  stron- 
tium carbonate  forms  strontium  oxide  and  carbon  dioxide. 

7.  Express  the  following  interactions  in  the  form  of  ionic  equations: 
(a)  calcium  chloride  and  sulphuric  acid;    (b)  strontium  nitrate  and  am- 
monium carbonate;   (c)  barium  acetate  and  potassium  chromate. 


CHAPTER   XXVII 


ALUMINIUM  —  CLAY  AND   CLAY  PRODUCTS 

429.  Occurrence.  —  Aluminium,   or    Aluminum,    does 
not  occur  free  in  nature,  but  its  compounds  are  numerous, 
abundant,  and  widely  distributed.     About  8  per  cent  of 
the  earth's  crust  is  combined  aluminium;  in  abundance 
it  ranks  first  among  the  metals  and   third  among  the 
elements  (8).     All  important  rocks  except  limestone  and 
sandstone  contain  silicates  of  aluminium  and  other  metals. 
Clay  and  slate  are  mainly  silicate  of  aluminium,  which 
was  formed  by  the  partial  decomposition  of  rocks  and 
minerals.     Corundum   and   emery  are   aluminium  oxide 
(A12O3).    Bauxite  is  an  hydroxide  of  aluminium  (H4A1205) ; 
it  is  often  colored  red  by  iron  oxide.     Cryolite  is  sodium 
aluminium  fluoride  (Na3AlF6);  it  is  a  white,  icelike  solid. 

430.  Metallurgy.  —  Aluminium    is    obtained     by    the 

electrolysis  of  alumin- 
ium oxide  (A^Os). 
The  purified  oxide  is 
dissolved  in  molten 
cryolite,  and  when  the 
current  passes,  alu- 
minium is  deposited 
at  the  cathode. 


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Fig.  87.  —  Apparatus  for  the  Manufacture 
of  Aluminium  by  the  Electrolysis  of 
Aluminium  Oxide.  Connection  is  Made 
with  the  Cathode  at  D. 


In  the  commercial  prep- 
aration of  aluminium,  an 
open  iron  vessel  (C,  C,  C)  lined  with  carbon  is  made  the  cathode 
(Fig.  87).  The  anode  consists  of  several  carbon  bars  (A,  A,  etc.)  hung 
from  a  common  copper  rod  (R),  which  can  be  lowered  as  the  carbon 


ALUMINIUM  329 

is  consumed.  The  bottom  of  the  box  is  first  covered  with  cryolite, 
the  anode  is  lowered,  and  the  box  is  then  filled  with  cryolite.  The  cur- 
rent is  turned  on,  and  in  its  resisted  passage  through  the  cryolite, 
enough  heat  is  generated  to  melt  the  cryolite.  Pure,  dry  aluminium 
oxide  is  now  added,  which  is  decomposed  into  aluminium  and  oxygen. 
The  oxygen  goes  to  the  anode,  where  part  escapes  and  part  unites 
with  the  carbon.  The  molten  aluminium,  being  heavier  than  cryolite, 
sinks  to  the  bottom  of  the  vessel.  The  process  is  continuous,  fresh 
aluminium  oxide  being  added  and  the  molten  aluminium  being  drawn 
off  at  intervals.  The  cryolite  is  unchanged  chemically. 

431.  Properties.  —  Aluminium  is  a  lustrous  white 
metal.  It  is  very  light  compared  with  other  common 
metals,  since  its  specific  gravity  is  only  about  2.6;  this 
value  is  one  third  that  of  iron.  It  is  ductile  and  malleable, 
and  is  extensively  made  into  wire  and  sheets.  It  is  a 
good  conductor  of  heat  and  electricity.  Compared  with 
most  common  metals  aluminium  is  rather  hard  and  strong. 
It  melts  at  about  658°  C.,  and  can  be  cast  and  welded, 
though  not  readily  soldered  so  as  to  produce  a  perma- 
nent joint.  Aluminium  is  only  very  slightly  acted  upon 
by  air  and  water.  It  combines  with  some  non-metals, 
especially  chlorine  and  sulphur;  it  combines  with  oxygen 
at  high  temperatures  and  is  an  excellent  reducing  agent 
(see  Thermit,  432).  Hydrochloric  acid  changes  it  into 
soluble  aluminium  chloride,  thus :  - 

2A1      +  6HC1  2A1C13  +     3H2 

Aluminium         Hydrochloric  Acid         Aluminium  Chloride          Hydrogen 

Under  ordinary  conditions  nitric  and  dilute  sulphuric 
acids  do  not  affect  it;  concentrated  sulphuric  acid  acts 
upon  it,  forming  aluminium  sulphate.  Sodium  chloride 
attacks  it,  especially  if  dilute  acids  are  present.  Sodium 
and  potassium  hydroxides  change  it  into  aluminates  with 
liberation  of  hydrogen,  thus :  - 


330  CHEMISTRY 

6NaOH         +      2A1  2Na3AlO3      +     3H2 

Sodium  Hydroxide        Aluminium        Sodium  Aluminate        Hydrogen 

432.  Uses. --The  varied  properties  of  aluminium, 
especially  its  strength,  lightness,  and  durability,  adapt 
it  to  numerous  uses.  It  is  made  into  the  metallic  parts 
of  military  outfits,  caps  for  jars,  surgical  instruments, 
cooking  utensils,  tubes,  fittings  of  boats,  automobiles, 
and  air  ships,  parts  of  opera  glasses  and  telescopes,  frame- 
work of  cameras,  stock  patterns  for  foundry  work,  hard- 
ware samples,  and  scientific  apparatus.  Its  attractive 
appearance  has  led  to  its  extensive  use  as  an  ornamental 
metal,  both  in  interior  decorative  work  and  in  numerous 
small  objects.  Aluminium  leaf  is  used  for  decorating 
book  covers  and  signs;  the  powder  suspended  in  an 
adhesive  liquid  is  used  as  a  protective  and  attractive 
paint  for  steam  pipes,  radiators,  smokestacks,  and  other 
metal  objects  exposed  to  heat  or  the  weather.  Alumin- 
ium wire  has  come  into  general  use  as  a  conductor  of 
electricity.  Large  quantities  of  aluminium  are  consumed 
in.  the  steel  industry,  in  the  manufacture  of  alloys  and 
certain  metals,  and  in  welding  metals. 

The  use  of  aluminium  in  the  steel  industry  depends  on 
the  ease  with  which  it  combines  with  any  oxygen  in  the 
molten  metal,  thereby  preventing  the  formation  of  small 
holes  in  the  solidified  metal.  Its  use  in  the  manufacture 
of  certain  metals  and  in  welding  is  based  on  its  property 
of  reducing  oxides  with  the  liberation  of  a  large  amount 
of  heat.  For  example,  when  a  mixture  of  chromium 
oxide  and  powdered  aluminium  is  ignited  at  one  point 
by  a  special  device,  the  reduction  thus  initiated  proceeds 
rapidly  throughout  the  mixture  and  the  intense  heat  fuses 
the  chromium,  which  can  be  removed  from  the  crucible 
subsequently  as  a  coherent  mass;  the  aluminium  oxide 


ALUMINIUM 


likewise  melts  and  separates  from  the  metal  as  a  slag. 
The  equation  for  the  chemical  change  is  - 

Cr2O3        +       2A1       =      2Cr     +         A12O3 

Chromium  Oxide        Aluminium        Chromium        Aluminium  Oxide 

Other  metals  hitherto  rare  or  expensive  are  similarly 
prepared.  If  a  mixture  of  fer- 
ric oxide  (Fe2O3)  and  powdered 
aluminium  is  ignited,  molten 
iron  at  a  temperature  of  about 
3000°  C.  is  produced.  An  ap- 
paratus has  been  devised  by 
which  the  molten  iron  can  be 
conducted  from  the  crucible 
into  a  mold  around  a  joint 
or  fracture  (Fig.  88).  By  this 
method  steel  rails  can  be 
quickly  welded  and  heavy  iron 
objects  repaired.  These  mix- 
tures of  aluminium  and  oxides 
are  called  thermit,  and  the  method  is  known  as  the 
Goldschmidt  or  aluminothermic  method.  (See  Part  II, 
Exp.  208.) 

433.  Alloys.  —  The  alloy  of  aluminium  and  copper '—  aluminium 
bronze  —  has  been  described  (399).     Magnalium  contains  from  75 
to  90  per  cent  of  aluminium,  the  rest  being  magnesium;  it  is  a  silvery, 
hard,  light  alloy,  which  takes  a  high  polish  and  is  very  durable.     It 
is  used  in  the  construction  of  chemical  balances  and  other  scientific 
instruments. 

434.  Aluminium  Oxide,  A12O3,  is    the    only  oxide    of 
aluminium.     It  is  often  called  alumina,  as  silicon  dioxide 
is  called  silica.     Its  native  forms,  corundum  and  emery, 
are  found  in  many  parts  of  the  United  States;    large 
quantities  of  emery  come  from  Asia  Minor  and  the  islands 


Fig.  88.—  Crucible  and  Mold  in 
Position  for  Welding  a  Steel 
Rail  with  Thermit. 


332  CHEMISTRY 

near  Greece.  Both  are  very  hard  substances,  pure 
corundum  ranking  next  to  diamond.  Emery  in  the  form 
of  powder,  cloth,  paper,  and  wheels  is  used  to  grind  and 
polish  hard  metals,  plate  glass,  etc.  The  transparent 
varieties  of  corundum  have  long  been  prized  as  gems, 
among  them  being  the  sapphire  and  ruby. 

Pure  aluminium  oxide  can  be  prepared  by  heating  its  hydroxide. 
Thus  prepared,  it  is  a  white  powder,  insoluble  in  water.  It  melts 
at  about  1900°  C.  Gems  having  the  same  composition  and  proper- 
ties as  the  natural  stones  are  now  made  from  aluminium  oxide  by 
melting  it,  alone  or  mixed  with  coloring  substances,  with  an  oxyhydro- 
gen  blowpipe.  If  more  or  less  pure  alumina  is  melted  in  an  electric 
furnace,  it  solidifies  on  cooling  to  a  crystalline  mass,  which  resembles 
corundum.  It  is  used  as  an  abrasive  and  resistent  material  under 
the  name  of  alundum.  When  alumina  or  any  other  compound  of 
aluminium  is  heated,  then  cooled  and  moistened  with  cobaltous 
nitrate  solution  and  heated  again,  the  mass  turns  a  beautiful  blue 
color.  Its  formation  is  a  test  for  aluminium. 

435.  Aluminium  Hydroxide,  A1(OH)3,  is  a  white, 
jellylike  solid  formed  by  adding  ammonium  hydroxide 
to  a  solution  of  an  aluminium  salt,  thus :  - 

AlCU     +  3NH4OH  =  A1(OH)S  +  3NH4C1 

Aluminium         Ammonium         Aluminium         Ammonium 
Chloride  Hydroxide  Hydroxide  Chloride 

It  is  insoluble  in  water.  It  interacts  with  strong  acids 
and  strong  bases  (in  excess),  forming  respectively  alumin- 
ium salts  and  aluminates.  Hence  it  may  be  regarded 
as  having  both  basic  and  acid  properties,  though  both 
are  weak.  Equations  illustrating  these  properties  are  - 

A1(OH)3          +      3HC1  A1C13         +  3H20 

Aluminium  Hydroxide    Hydrochloric  Acid     Aluminium  Chloride       Water 

A1(OH)3  +  3NaOH   =  Na3AlO3  +  3H2O 

Aluminium  Sodium  Sodium  Water 

Hydroxide  Hydroxide        Aluminate 


ALUMINIUM  333 

436.  Aluminium  Sulphate,  A12(S04)3,  is  a  white  solid 
prepared  from  clay  or  bauxite  by  heating  with  sulphuric 
acid.     The  commercial  substance  is  often  impure.     The 
crystallized  salt  usually  has  the  formula  Al2(SO4)3.i8H2O. 
It  is  used  in  dyeing  and  paper  making,  in  purifying  water, 
and  in  preparing  alum  and  other  aluminium  compounds. 

A  solution  of  aluminium  sulphate  has  an  acid  reaction  on  account 
of  hydrolysis ;  the  equation  for  the  hydrolysis  is  — 

Al2(S04)3       +  6H20  =         2A1(OH)3        +    3H2S04 

Aluminium  Sulphate          Water          Aluminium  Hydroxide         Sulphuric  Acid 

The  acid  reaction  is  due  to  the  hydrogen  ions  produced  by  the  ioniza- 
tion  of  the  sulphuric  acid.  Practical  application  is  made  of  this 
reaction  in  purifying  water.  Upon  adding  aluminium  sulphate  to 
impure  water,  the  gelatinous  aluminium  hydroxide  that  is  precipi- 
tated slowly  settles  and  carries  with  it  suspended  particles  and  germs. 
(See  Part  II,  Exp.  207.) 

437.  Alum.  —  When  solutions  of   aluminium  sulphate 
and  potassium  sulphate  are  mixed  and  concentrated  by 
evaporation,    transparent,    colorless,    glassy   crystals   are 
deposited.     This  solid  is  potassium  alum  or  simply  alum. 
It    has    the    composition    represented    by    the    formula, 
K2A12(SO4)4.24H2O,  or  K2SO4.A12(SO4)3.24H20.     It  is  the 
type  of  a  class  of  similar  salts  called  alums,  which  can 
be  prepared  from  sulphates  of  univalent  and  trivalent 
metals  (e.g.  K,  Na,  NH4,  and  Al,  Cr,  Fe). 

Alums  are  rather  soluble  in  water,  and  their  solutions 
have  an  acid  reaction  owing  to  hydrolysis  (compare  436) . 
They  crystallize  as  octahedrons  and  contain  twenty-four 
molecules  of  water  of  crystallization.  When  heated,  alums 
lose  their  water  of  crystallization  and  usually  some  sul- 
phur trioxide,  and  become  a  white  powder  or  a  porous 
mass  known  as  burnt  alum. 

Potassium  alum  is  the  most  common,  but  ammonium 


334  CHEMISTRY 

and  sodium  alums  are  manufactured  and  used.  Alum 
(and  sometimes  aluminium  sulphate)  is  an  ingredient 
of  some  baking  powders;  such  powders  are  not  to  be 
recommended  because  they  interfere  with  digestion. 
Alums  are  used  in  dyeing  and  printing  cloth,  in  tan- 
ning and  paper  making,  as  a  medicine,  for  hardening 
plaster,  in  making  wood  and  cloth  fire-proof,  and  in 
preparing  aluminium  compounds.  Aluminium  sulphate 
is  gradually  displacing  aluminium  alums  for  many  pur- 
poses, especially  the  purification  of  water.  (See  Part  II, 
Exps.  207,  210,  211.) 

438.  Mordants.  —  Aluminium   hydroxide   is   extensively  used   as 
a  mordant  in  dyeing.     Many  dyes  must  be  fixed  in  the  fiber  by  a 
metallic  substance,  otherwise  the  color  would  be  easily  removed. 
The  cloth  to  be  dyed  or  printed  is  first  impregnated  or  printed  with 
an  aluminium  salt,  such  as  aluminium  acetate,  and  then  exposed  to 
steam  or  treated  with  ammonium  hydroxide.     This  operation  changes 
the  aluminium  salt  into  aluminium  hydroxide,  which  is  precipitated 
in  the  fiber  of  the  cloth.     The  mordanted  cloth  is  next  passed  through 
a  vat  containing  the  solution  of  the  dye,  which  unites  chemically 
or  mechanically  (perhaps  both)  with  the  aluminium  hydroxide,  form- 
ing a  colored  compound.     The  latter  is  relatively  insoluble  and  cannot 
be  easily  washed  from  the  cloth,  i.e.  it  is  a  fast  color. 

439.  Aluminium  Chloride,  A1C13,  when  pure  is  a  white   powder, 
but  it  is  often  a  yellowish,   crystalline  mass   (A1C13.6H2O).     It  is 
prepared  by  heating  powdered  aluminium  in  chlorine,  or  by  passing 
chlorine  over  a  heated  mixture  of  aluminium  oxide  and  carbon. 
Exposed  to  air,  it  absorbs  moisture  and  gives  off  fumes  of  hydro- 
chloric acid.     It  dissolves  in  water  with  evolution  of  heat,  and  if 
the  solution  is  heated,  hydrochloric  acid  is  expelled,  owing  to  the 
hydrolysis  of  the  chloride,  thus  — 


A1C13       +    3H20  =      3HC1      +  A1(OH), 

Aluminium  Water         Hydrochloric  Aluminium 

Chloride  Acid  Hydroxide 


CLAY  AND  CLAY  PRODUCTS       335 

Clay  and  Clay  Products 

440.  Clay  is  a  more  or  less  impure  aluminium  silicate, 
formed  by  the  slow  decomposition  of  rocks  containing 
aluminium    compounds,    especially   the   feldspars.     Pure 
feldspar  is  a  silicate  of  aluminium  and  sodium  or  potas- 
sium   (NaAlSi308   or   KAlSi3O8).      The   products   of    its 
decomposition  are  chiefly  an  insoluble  aluminium  silicate 
and   a  soluble   alkaline   silicate.     The  latter  is  washed 
away.     The  pure   aluminium   silicate   which   remains  is 
called  kaolin  (H4Al2Si2O9) .     Usually  kaolin  is  mixed  with 
particles  of  mica  and  quartz,  carbonates  of   calcium  and 
magnesium,  and  iron  compounds.     This  mixture,  which 
varies  in  composition,   is  known  as  clay.     Kaolin  is  a 
white,  powdery  solid.     It  becomes  slightly  plastic  when 
wet,  and  can  therefore  be  molded  into  various  shapes. 
Ordinary  clay  is  very  plastic  when  wet,  more  easily  fused 
than  kaolin,  but  shrinks  when  dried  and  burned;   it  also 
contains   iron    compounds,    which    color   it   gray,    blue, 
yellow,  brown,  and  red.     All  clays  have  a  peculiar  clayey 
odor  when  moist.     If  heated,  clay  and  clay  mixtures  do 
not  melt  (except  at  a  very  high  temperature),  but  upon 
cooling  become  very  hard. 

441.  Clay  Products.  —  Porcelain  is  made  by  mixing  kaolin,  fine 
sand,  and  powdered  feldspar,  shaping  the  mass  into  the  desired  form 
by  molds  or  on  a  potter's  wheel,  and  then  heating  in  a  kiln  to  a 
high  temperature.     The  mass  when  cool  is  hard  and  translucent  (if 
thin);  it  is  not  easily  corroded  by  chemicals  (except  fused  alkalies). 
Although  it  is  not  very  porous,  its  surface  is  glazed,  partly  for  pro- 
tection, partly  for  ornament.     This  is  done  by  coating  it  with  a  mix- 
ture similar  to  that  used  for  making  the  porcelain  but  more  easily 
fused,  and  then  heating  again  so  that  the  glaze  will  melt  and  penetrate 
the  surface.     Porcelain  is  often  decorated  by  painting  or  printing 
designs  on  the  surface  with  metallic  paints  or  colored  glass  and  then 


336  CHEMISTRY 

heating  again.  In  making  pottery  the  raw  materials  are  less  care- 
fully selected  and  prepared,  and  not  heated  to  such  a  high  temper- 
ature. The  best  grades  can  hardly  be  distinguished  from  porcelain, 
but  usually  pottery  is  much  heavier  and  thicker.  If  less  pure,  plastic 
clay  is  used  and  heated  to  a  moderate  temperature,  the  product  is 
known  as  earthenware.  This  is  a  large  class  and  includes  tiles,  terra- 
cotta, jugs,  flowerpots,  and  clay  tobacco  pipes.  This  ware  is  porous 
and  is  sometimes  glazed  by  throwing  salt  into  the  kiln  just  before  the 
operation  is  over.  The  salt  volatilizes  and  forms  a  fusible  sodium 
aluminium  silicate  upon  the  surface.  Clay  products  used  for  con- 
struction include  bricks,  conduits,  drain  pipe,  etc.  They  are  made 
from  impure  clay  and  heated  just  enough  to  harden  the  mixture. 
The  product  varies  with  the  clay,  but  is  often  colored  red  owing  to 
iron  oxide  formed  from  the  iron  compounds  in  the  unburned  clay. 

EXERCISES 

1.  Compare  the  proportion  of  aluminium  in  the  earth's  crust  with  the 
proportion  of  other  abundant  elements  (8). 

2.  Compare  the  metallurgy  of  aluminium  with  that  of  other  metals. 

3.  Discuss  the  interaction  of  aluminium  hydroxide  with  acids  and  with 
alkalies. 

4.  Essay  topics:    (a)  History  of  aluminium,     (b)  Ceramics,     (c)  Clay. 
(d)  Dyeing,     (e)  Uses  of  aluminium.     (/)  Aluminium  in  gems,     (g)  Use  and 
care  of  aluminium  ware. 

6.  Practical  topics:  (a)  Two  pans  are  identical  in  size  and  thickness  of 
material,  one  being  of  aluminium  and'  the  other  of  iron.  Compare  their 
weights,  (b)  Starting  with  aluminium  how  would  you  prepare  in  succes- 
sion AlCls,  A1(OH)3,  Na3AlO3,  A1C13,  A1(OH)3,  A12O3,  Al  ? 

PROBLEMS 

1.  Calculate  the  weight  of   aluminium  in   (a)  20   gm.   of   aluminium 
oxide,  (b)  34  gm.  of  aluminium  hydroxide,  (c)  49  gm.  of  aluminium  chloride. 

2.  Honeystone  is  a  mineral  occurring  in  brown  coal  seams  and  has  the 
formula  Al3Ci2Oi2.i8H2O.     Calculate  the  per  cent  of  aluminium  in  it. 

3.  Write  the  formulas  of  the  following  compounds:  Aluminium  nitrate, 
aluminium   bromide,   aluminium  phosphate    (ortho),   aluminium   fluoride, 
aluminium  silicate  (meta),  potassium  aluminate.     Calculate  the  per  cent  of 
aluminium  in  any  three  of  these  compounds. 

4.  How  many  pounds  of  aluminium  in  a  ton  (2000  Ib.)  of  pure  kaolin? 

5.  What  weight  of  aluminium  can  be  obtained  from  100  kilograms  of 
bauxite  (93  per  cent  A1(OH)3)? 


CHAPTER  XXVIII 
IRON  — NICKEL  AND   COBALT 

442.  Occurrence  of  Iron.  —  In   abundance  iron  ranks 
fourth    among    the    elements    and    second    among    the 
metals    (8).     Uncombined  iron  is   found  in   meteorites. 
Combined  iron. is  found  in  most  rocks,  soils,  and  natural 
waters.     It  is  assimilated  by  plants  and  animals  and  is 
essential  to  their  life  processes,  being  a  constituent  of 
chlorophyll  (the  green  coloring  matter  of  plants)  and  of 
hemoglobin  (the  red  coloring  matter  of  blood).     The  chief 
ores  of  iron  are  red  and  black  hematite  (Fe2O3),  brown 
hematite  (limonite,  Fe4O3(OH)6),  magnetite  (FegC^),  and 
siderite  (FeCOs).     The  most  abundant  ore  and  the  chief 
source  of  iron  and  steel  is  hematite,  which  comes  mainly 
from  the  Lake  Superior  region.  Large  quantities  of  iron  ore 
are  also  mined  in  Alabama,  Tennessee,  and  the  Virginias. 

Other  abundant  compounds  of  iron,  not  used  as  a  source  of  the 
metal,  are  iron  pyrites  (FeS2),  pyrrhotite  (varying  from  Fe6S7  to 
FenSi2),  and  the  copper-iron  sulphides  (chalcopyrite,  CuFeS2,  and 
bornite,  Cu3FeS3  —  see  394). 

443.  Preparation     and    Properties    of    Pure    Iron.  — 

Chemically  pure  iron,  though  uncommon  in  commerce, 
may  be  obtained  as  a  powder  by  reducing  the  oxide  with 
hydrogen  or  as  irregular  plates  by  the  electrolysis  of  a 
solution  of  ferrous  sulphate.  Such  iron  is  called  iron 
"by  hydrogen,"  or  electrolytic  iron. 

Pure  iron  is  a  silvery  white,  lustrous  metal.  It  is 
ductile  and  malleable,  and  softer  than  ordinary  iron, 
being  about  as  soft  as  aluminium.  The  specific  gravity 


338  CHEMISTRY 

is  7.86  and  the  melting  point  is  1520°  C.  It  is  attracted 
by  a  magnet,  but  soon  loses  its  own  magnetism.  Dry 
air  has  no  effect  upon  iron,  but  moist  air  rusts  it.  Rust- 
ing is  a  complex  process  and  is  explained  in  different 
ways.  A  recent  and  acceptable  interpretation  based  on 
the  theory  of  electrolytic  dissociation  is  that  the  iron 
first  interacts  with  the  water.  The  iron  goes  into  solu- 
tion as  ferrous  ions  (Fe++)  and  hydrogen  ions  (H+) 
escape  as  hydrogen  gas  (H2);  the  ferrous  ions  combine 
with  the  hydroxyl  ions  (OH~)  left  in  the.  water  and  form 
ferrous  hydroxide  (Fe(OH)2),  which  is  subsequently  con- 
verted into  the  complex  substance  called  iron  rust.  Once 
begun,  rusting  proceeds  rapidly,  because  the  film  of  rust 
is  not  compact  enough  to  protect  the  metal.  Like  many 
metals,  iron  readily  interacts  with  dilute  acids,  and  as  a 
rule  hydrogen  and  ferrous  compounds  (e.g.  ferrous  sul- 
phate, FeS04)  are  the  products. 

444.  Varieties    of   Iron.  —  It    is   customary   to   speak 
of  three  varieties  of  iron,  —  cast  iron,  wrought  iron,  and 
steel.     The  physical  properties  of  the  different  varieties 
of  iron  are  modified  by  the  proportions  of  carbon  and 
the   other   constituents   as    well    as    by  the    method   of 
manufacture.      Hence  there  are  several  different  kinds 
of  iron  and  steel. 

445.  Metallurgy    of    Cast    Iron.  —  Iron    is    extracted 
most  easily  from  its  oxides.     The  ore,  if  not  an  oxide,  is 
first  crushed  and  roasted  to  change  it  into  ferric  oxide 
(Fe2Os)   as  far  as  possible.     Thus  prepared,   the  ore  is 
mixed  with  a  flux  (usually  limestone)  and  carbon  (usually 
in  the  form  of  coke)   and  smelted  in  a  blast  furnace. 
The  carbon  reduces  the  oxide  to  metallic  iron  and  the 
flux    converts    impurities    in    the  ore   (e.g.    silicon    and 
aluminium  compounds)  into  fusible  silicates  called  slag. 


IRON  —  NICKEL  AND   COBALT 


339 


A  blast  furnace  (Fig.  89)  is  a  huge  tower,  about  ninety  feet  high 
and  twenty  feet  in  diameter 
at  the  largest  part;  it  is  nar- 
rower at  the  top  and  bottom 
than  in  the  middle.  It  is 
built  of  steel  and  lined  with 
fire  brick.  There  are  pipes  at 
the  bottom,  called  tuyeres, 
through  which  large  quanti- 
ties of  hot,  dry  air  are  forced 
into  the  furnace,  thereby  pro- 
ducing the  high  temperature 
required  in  the  smelting;  while 
another  pipe  at  the  top  not 
only  permits  the  escape  of  hot 
gaseous  products,  but  con- 
ducts them  into  a  series  of 
pipes  which  lead  to  different 
parts  of  the  plant,  where  the 
hot  gases  are  utilized  to  heat 
the  air  blown  through  the  tu- 
yeres, and  also  as  fuel. 

When  the  blast  fur- 
nace has  been  heated  to 
the  proper  temperature, 
or  is  already  in  opera- 
tion, charges  of  the 
proper  proportions  of 
ore,  coke,  and  flux  are 
introduced  at  intervals 
into  the  furnace  by 
dumping  them  upon  the 
cone-shaped  cover.  As 
the  hot  air  enters  at  the 


Fig.  89.  —  Blast  Furnace.  A,  Throat; 
B,  Bosh;  C,  Crucible  Where  the  Melted 
Iron  Collects;  D,  Pipe  for  Hot  Air 
Blast  to  the  Tuyeres  (T);  E,  Escape 
Pipe  for  Gases  Which  do  not  Escape 
Through  the  " Down  Comer";  G,  Cup; 
H,  Cone;  N,  Trough  for  Drawing  off 
Slag;  T,  Tuyere;  I,  Hole  Through 
Which  Iron  is  Withdrawn. 


bottom,  a  portion  of  the   carbon   forms   carbon  dioxide, 
which  is  reduced  by  hot  carbon  to  carbon  monoxide.    The 


340  CHEMISTRY 

latter  interacts  with  the  ore  and  reduces  it.  As  the 
smelting  proceeds,  the  reduction  continues  and  the  flux 
forms  a  slag;  both  iron  and  slag  sink,  the  molten  iron 
finally  falling  through  the  slag  to  the  bottom  of  the 
furnace,  where  both  are  drawn  off  through  separate 
openings  at  desired  intervals.  The  iron  is  usually  run 
from  the  furnace  into  molds  of  sand  or  iron  and  allowed 
to  solidify;  such  iron  is  often  called  pig  iron.  In  some 
plants  the  molten  iron  is  run  into  huge  vessels,  called 
converters,  and  made  directly  into  steel  (see  449  (i)). 

446.  Properties  of  Cast  Iron.  —  Cast  iron  contains  from 
2  to  5  per  cent  of  carbon,  together  with  varying  propor- 
tions of  silicon  and  manganese  and  traces  of  phosphorus 
and  sulphur.  If  most  of  the  carbon  is  combined  with  the 
iron,  the  variety  is  called  white  cast  iron.  But  if  much 
of  the  carbon  remains  uncombined  as  graphite,  gray  cast 
iron  is  formed;  it  is  softer  and  less  brittle  than  the  white 
variety,  and  melts  at  a  lower  temperature.  Cast  iron  has 
a  crystalline  structure  and  is  brittle;  it  will  withstand 
great  pressure.  It  cannot  be  welded  or  forged,  that  is,  hot 
pieces  cannot  be  united,  nor  be  shaped  by  hammering. 
But  it  can  be  cast.  Cast  iron  is 'not  attacked  by  alkalies 
and  only  slightly  by  concentrated  acids.  Sulphuric  acid 
is  sometimes  concentrated  and  often  transported  in  iron 
vessels.  Dilute  acids  interact  readily  with  cast  iron. 

This  is  the  kind  of  iron  used  in  an  ordinary  iron  foun- 
dry. The  iron  which  melts  at  a  comparatively  low  tem- 
perature (about  1200°  C.)  is  heated  in  a  furnace  similar 
to  a  blast  furnace,  and  when  molten  is  poured  into  sand 
molds  of  the  desired  shape.  Considerable  cast  iron  is 
made  into  steel.  Cast  iron  containing  5  to  20  per  cent 
of  manganese  is  called  spiegel  iron,  while  ferro-manganese 
contains  20  or  more  per  cent  of  manganese. 


IRON  — NICKEL   AND    COBALT 


447.    Preparation  and  Properties  of  Wrought  Iron.  - 
Wrought  iron  is  made  from  cast  iron  by  removing  most 
of  the  impurities  (carbon,  silicon,  phosphorus,  and  sul- 
phur). 

The  process  is  conducted  in  a  furnace  much  like  a 
reverberatory  furnace  (Fig.  90).  The  bottom  (B)  of  the 
furnace  is  covered  with 
iron  ore  (ferric  oxide, 
Fe2O3)  and  the  charge 
of  cast  iron  and  flux  is 
laid  upon  it.  The  in- 
tense heat  that  is  re- 
flected down  upon  the 
charge  by  the  roof  of 
the  furnace  melts  the 
cast  iron;  the  carbon 


Fig.  90.  —  Reverberatory  Furnace.  The 
fire  burns  on  the  grate,  G,  and  the 
long  flame  which  passes  over  the  bridge, 
E,  is  reflected  down  by  the  sloping  roof 
upon  the  contents  of  the  furnace. 
Gases  escape  through  I.  The  charge, 
which  rests  upon  B,  does  not  come  in 
contact  with  the  fuel,  but  is  oxidized 
or  reduced  by  the  flame. 


forms  carbon  monoxide 
and  dioxide  which  es- 
cape, while  the  silicon, 
manganese,  sulphur, 
and  phosphorus  are  oxi- 
dized and  unite  with  the  flux  or  the  iron  oxide  and  form 
a  slag.  As  the  impurities  are  removed,  the  mass  becomes 
pasty,  owing  to  its  higher  melting  point,  and  is  stirred 
vigorously,  or  " puddled."  At  the  proper  time  large 
lumps  called  blooms  are  removed  and  hammered,  or  more 
often  rolled  between  ponderous  rollers.  This  operation 
removes  most  of  the  slag,  and  if  the  rolling  is  repeated, 
the  quality  of  the  iron  is  improved;  the  final  rolling  often 
leaves  the  iron  in  the  desired  shape. 

Wrought  iron  is  the  purest  variety  of  commercial  iron. 
It  is  practically  pure  iron  mixed  with  a  little  slag  (.1  to 
2  per  cent).  The  iron  itself  contains  not  more  than  0.5 


342  CHEMISTRY 

per  cent  of  carbon  and  sometimes  only  0.06  per  cent, 
the  average  being  about  0.15  per  cent;  the  other  elements 
are  present  in  mere  traces.  Unlike  cast  iron  it  is  fibrous 
and  can  be  bent.  It  melts  at  a  high  temperature  (1600° 
to  2000°  C.)  and  it  is  not  used  for  casting.  Since  it 
softens  at  about  1000°  C.,  it  can  be  forged  and  welded, 
and  may  be  seen  undergoing  these  operations  in  a  black- 
smith's shop.  It  is  very  malleable  and  ductile,  and  in 
these  forms  the  metal  is  very  strong.  It  can  be  rolled 
into  plates  and  sheets  and  drawn  into  fine  wire.  Wrought 
iron  rusts  more  rapidly  than  cast  iron,  and  is  also  more 
vigorously  attacked  by  acids  and  alkalies  at  a  high  tem- 
perature. Wrought  iron  is  made  into  wire,  sheets,  rods, 
nails, .  spikes,  bolts,  chains,  anchors,  horseshoes,  tires, 
and  agricultural  implements.  It  is  less  important  than 
formerly,  since  it  is  being  replaced  by  soft  steel  (452). 

448.  Steel  is  a  form  of  iron  which  contains,  as  a  rule, 
more  carbon  and  other  elements  than  wrought  iron  and 
less  than  cast  iron.     However,  many  grades  of  steel  are 
manufactured,    and    their    physical    properties    do    not 
depend  merely  on  the   presence   of   certain  proportions 
of  carbon,  phosphorus,  silicon,  sulphur,  etc.,  but  more 
especially  on  the  method  of  manufacture. 

449.  Manufacture  of  Steel.  —  Steel  is  made  from  cast 
iron.     The  aim  in  the  manufacture  is  to  prepare  a  pro- 
duct containing  little   or  no   sulphur,   phosphorus,   and 
silicon,  but  the  desired  proportion  of   carbon.     Most  of 
the  steel  is  made  in  the  United  States  by  two  general 
methods.     The  Bessemer  process  consists  in  burning  out 
the  impurities  in  cast  iron  by  forcing  air  through  the 
molten  metal  contained  in  a  movable  receptacle  lined 
with  silica,  and  then  adding  just  enough  iron  of  known 
composition  to  purify  the  metal  and  give   the  desired 


IRON  — NICKEL    AND    COBALT 


343 


proportion  of  carbon.  In  the  open-hearth  process  a  mix- 
ture of  cast  iron  and  other  iron  products  is  heated  by 
hot  gases  on  the  hearth  or  bed  of  a  furnace  lined  with 
limestone  or  dolomite. 

(i)  The  Bessemer  process  is  carried  on  in  a  converter 
(Fig.  91).     This  is  a  huge,  pear-shaped  vessel,  supported 


Air- 


Fig.  91.  —  Converter. 

so  that  it  can  be  tipped  into  different  positions;  it  is  also 
provided  with  one  hollow  trunnion  and  holes  (C,  C,  C)  at 
the  bottom,  through  which  a  powerful  blast  of  air  can 
be  blown.  It  is  made  of  thick  wrought-iron  plates  and 
is  lined  with  an  infusible  mixture  rich  in  silica.  The 
converter  when  in  use  is  swung  into  a  horizontal  posi- 
tion, and  fifteen  to  twenty  tons  of  molten  cast  iron  are 
run  in,  often  directly  from  the  blast  furnace.  The  air 
blast  is  turned  on,  and  the  converter  is  swung  back 
to  a  vertical  position.  As  the  air  is  forced  in  fine  jets 
through  the  molten  metal,  the  temperature  rises,  and  the 
carbon,  silicon,  and  manganese  are  oxidized.  The  carbon 
forms  carbon  monoxide,  which  burns  at  the  mouth  of  the 
converter  in  a  large  brilliant  flame,  while  the  other  oxides 
pass  into  the  slag.  This  oxidation  generates  enough 
heat  to  keep  the  metal  melted.  As  soon  as  the  impuri- 


344  CHEMISTRY 

ties  have  been  burned  out,  sufficient  spiegel  iron  or  ferro- 
manganese  is  added  to  furnish  the  proper  amount  of 
carbon  and  manganese.  By  adding  certain  metals,  e.g. 
aluminium,  titanium,  or  vanadium,  gases  are  removed, 
and  a  better  quality  of  steel  is  produced.  After  the 
completion  of  the  operation,  which  takes  about  twenty 
minutes,  the  converter  is  tilted  and  the  metal  is  poured 
into  molds  to  form  blocks  called  ingots,  which  are  subse- 
quently shaped  into  rails  or  other  objects. 

The  process  described  in  the  preceding  paragraph  is  called  the  acid 
Bessemer  process  because  the  converter  is  lined  with  silica,  which  is 
a  non-metallic  or  an  acid  anhydride  (187).  By  this  process  the  carbon 
and  silicon  can  be  removed  but  not  all  the  sulphur  and  phosphorus. 
It  is  used  in  the  United  States  because  domestic  ores  are  low  in 
phosphorus  and  sulphur.  In  Europe  the  Thomas-Gilchrist  process  or 
basic  process  is  used.  The  converter  in  this  modified  process  is  lined 
with  burned  dolomite  (i.e.  practically  a  mixture  of  lime  and  magnesia). 
This  lining,  after  use,  is  known  as  Thomas  slag;  it  is  utilized  as  a 
fertilizer  on  account  of  its  phosphorus  content. 

(2)  The  open-hearth  process  is  conducted  in  a  special 
kind  of  furnace  called  an  open-hearth  furnace  (Fig.  92). 
The  receptacle,  or  hearth  (H),  in  which  the  charge  is 
melted,  is  lined  with  limestone  or  burned  dolomite, 
thereby  making  the  process  a  basic  one.  'At  the  base  of 
the  furnace  are  duplicate  sets  of  checkerwork  (A,  B  and 
C,  D)  arranged  for  alternate  use.  As  the  hot  gases  pass 
through  A,  B  to  the  chimney,  they  heat  the  checkerwork. 
The  fuel  gas  is  then  passed  through  B  and  air  through  A. 
The  mixture  of  air  and  gas  burns  and  produces  a  much 
higher  temperature  on  the  hearth  than  if  the  gaseous 
mixture  were  cool.  The  oxidizing  flame  passes  over  the 
charge  on  the  hearth  (H),  oxidizing  some  of  the  impuri- 
ties and  keeping  the  mass  at  such  a  temperature  that 
other  impurities  form  a  slag  with  the  lining.  Meanwhile 


IRON  — NICKEL   AND   COBALT 


345 


the  hot  products  of  combustion  and  the  unused  gases 
are  passed  through  the  checkerwork  C,  D  and  heat  it. 
By  means  of  valves  the  fuel  gas  and  air  are  then  made 
to  pass  through  this  checkerwork  to  the  hearth  and 
out  over  the  other  checkerwork  (A,  B)  to  the  chimney. 


Fig.  92.  — •  Open-Hearth  Furnace.     (Vertical  Section.) 

Thus  the  process  is  alternated,  one  checkerwork  being 
cooled  as  the  other  is  heated,  and  vice  versa.  It  is  only 
by  this  regenerative  process,  as  it  is  called,  that  enough 
heat  is  obtained  to  keep  the  charge  melted  as  it  becomes 
purer  and  purer.  The  charge  is  heated  from  six  to 
twelve  hours;  when  a  test  shows  that  the  metal  contains 
the  desired  proportion  of  carbon,  ferromanganese  is 
added,  and  the  steel  is  quickly  poured  into  molds  and 
allowed  to  cool  into  ingots.  The  open-hearth  process 
is  easily  controlled  and  yields  a  tough,  elastic  steel, 
which  is  excellent  for  bridges,  large  machines,  large  guns, 
and  gun  carriages. 

450.  Other  Processes  of  Making  Steel.  —  The  crucible  process 

consists  in  melting  wrought  iron  with  charcoal  in  graphite  or  clay 


346  CHEMISTRY 

crucibles.  During  the  melting  the  iron  is  slowly  changed  into  steel 
by  absorbing  the  proper  proportions  of  carbon.  In  the  cementation 
process  wrought  iron  and  carbon  are  heated  in  fire-brick  boxes  for 
several  days.  The  transformation  is  the  same  as  in  the  crucible 
process.  In  the  electric  process  a  superior  quality  of  steel  is  made 
by  utilizing  the  heat  of  an  electric  furnace. 

451.  Alloys  of  Steel.  —  Steel  alloys,  or  special  steels 
as  they  are  called,  are  made  by  adding  to  steel  small 
quantities  of  certain  metals,  such  as  nickel,   chromium, 
molybdenum,     tungsten,     vanadium,     and     manganese. 
Such  steel  alloys  differ  in  special  properties,  though  all 
are  characterized  by  extreme  hardness,  toughness,  and 
strength,    which    make    them    almost   indispensable    for 
certain  uses.     Thus,   nickel-chromium  steel  is  used  for 
armor  plate,  tungsten  and  molybdenum  steel  for  high- 
speed tools,  manganese  steel  for  safes  and  rock-crushing 
machinery,    and   vanadium   steel   for   automobile   parts. 
The  metals  are  sometimes  added  directly  to  the  molten 
steel,  but  more  often  in  the  state  of  an  alloy  of  iron; 
these  alloys  contain  varying  but  known  percentages  of 
the  metals,  and  are  called  ferrochrome,  ferrosilicon,  etc. 

452.  Properties   of  Commercial  Steel. --The  proper- 
ties are  numerous  because  there  are  many  kinds  of  steel. 
Thus,    steel,    using    this    term    broadly,    is    fusible    and 
malleable,  and  can  be  forged,  welded,  and  cast.     Varie- 
ties  containing  0.2   per   cent  of   carbon   are  much  like 
wrought  iron  and  are  called  soft  or  mild  steel.    Structural 
steel  contains  more  carbon  (0.2  to  0.8  per  cent)  and  is 
hard  like  cast  iron,  while  tool  steel,  which  contains  up- 
wards of  1.5  per  cent  of  carbon,  is  very  hard  indeed.     Its 
most  valuable  property  is  the  varying  hardness  that  it 
may  be  made  to  acquire.     If  steel  is  heated  very  hot 
and  then  suddenly  cooled  by  immersion  in  cold  water 


IRON  — NICKEL  AND   COBALT  347 

or  oil,  it  becomes  brittle  and  very  hard.  But  if  heated 
and  cooled  slowly,  it  becomes  soft,  tough,  and  elastic. 
All  grades  of  hardness  may  be  obtained  between  these 
extremes.  And  if  the  hardened  steel  is  reheated  to  a 
definite  temperature,  determined  approximately  by  the 
color  the  oxidized  metal  assumes,  and  then  properly 
cooled,  a  definite  degree  of  hardness  and  elasticity  is 
obtained.  This  last  operation  is  called  tempering. 

453.  Uses  of  Steel.  —  Steel  is  now  used  instead    of 
iron   for    countless   purposes.     High   buildings,    bridges, 
rails,  cars,  locomotives,  battleships,  electrical  machinery, 
boilers,  agricultural  implements,  wire  nails,  rods,  hoops, 
tin  plates,  and  castings  of  all  kinds  consume  vast  amounts 
of   Bessemer   and   open-hearth   steel.     Crucible   steel   is 
used  in  making  springs,  tools,  cutlery,  pens,  and  needles. 

454.  Compounds  of  Iron.  —  Iron  forms  two  series  of 
compounds  —  ferrous  and  ferric.     They  are  analogous  to 
cuprous  and  cupric,  mercurous  and  mercuric  compounds. 
The  valence  of  iron  is  two  in  ferrous  compounds  and  three 
in  ferric.     Ferrous  compounds  pass  into  the  correspond- 
ing ferric  compounds  by  oxidation,  while  ferric  compounds 
become   ferrous   by   reduction.     The   passage   from   one 
series  to  the  other  occurs  easily,  especially  from  ferrous 
to  ferric.     Iron  also  forms  many  complex   compounds. 
(See  Part  II,  Exps.  217,  218,  219.) 

455.  Oxides.  —  Iron  forms  three  oxides.    Ferrous  oxide 
(FeO)    is     an    unstable    black    powder.      Ferric    oxide 
(Fe2O3)  occurs  native  as  hematite.     Large  quantities  are 
manufactured  by  heating  the  ferrous  sulphate  obtained 
as  a  by-product  in  the  cleaning  of  iron  castings,  rods, 
and  sheets.     It  is  sold  under  the  names  rouge,   crocus, 
and    Venetian    red,    and    is    used    to   polish    glass    and 
jewelry  and  to  make  red  paint.     Ferrous-ferric  or  fer- 


348  CHEMISTRY 

roso-ferric  oxide  (magnetic  oxide  of  iron,  Fe304)  occurs 
native  as  magnetite;  if  noticeably  magnetic,  it  is  called 
loadstone.  It  is  produced  as  a  black  film  or  scale  by 
heating  iron  in  the  air;  heaps  of  it  are  often  seen  beside 
the  anvil  in  a  blacksmith's  shop.  The  firm  coating  of 
this  oxide  formed  by  exposing  iron  to  steam  protects  the 
metal  from  further  oxidation;  iron  thus  coated  is  called 
Russia  iron.  Some  authorities  regard  this  oxide  as  iron 
ferrite  (Fe(FeO2)2). 

456.  Hydroxides.  —  Ferrous    hydroxide    (Fe(OH)i)    is    a    white 
solid  formed  by  the  interaction  of  a  ferrous  salt  and  an  hydroxide, 
such  as  sodium  hydroxide.     Exposed  to  the  air,  it  soon  turns  green, 
and  finally  brown,  owing  to  oxidation  to  ferric  hydroxide.     Ferric 
hydroxide  (Fe(OH)3)  is  a  reddish  brown  solid,  formed  by  the  inter- 
action of  an  hydroxide  and  a  ferric  salt. 

457.  Ferrous  Sulphate  (FeS04)  is  a  green  salt  obtained 
by  the  interaction  of  iron  (or  of  ferrous  sulphide)  and 
dilute  sulphuric  acid   (e.g.  see  455).     It  is  prepared  on 
a  large  scale  by  oxidizing  iron  pyrites   (FeS2).     This  is 
accomplished    simply    by    roasting,    or    more    often    by 
exposing  heaps  of  pyrites  to  moist  air;    the  mass  is  ex- 
tracted with  water  containing'  scrap  iron  and  a  small 
proportion  of  sulphuric  acid.     From  the  clear  solution, 
large   light   green   crystals   are   obtained    (FeS04.7H20); 
in  this  form  the  substance  is  often  called  green  vitriol 
or  copperas.     Exposed  to  the  air,  ferrous  sulphate  efflo- 
resces and  oxidizes.    Large  quantities  are  used  in  dyeing 
silk  and  wool,  as  a  disinfectant,  and  in  manufacturing 
ink,  bluing,  and  pigments.     Much  black  writing  ink  is 
made   essentially  by  mixing  ferrous  sulphate,   nutgalls, 
gum,  and  water. 

458.  Ferric  Sulphate   (Fe2(SO4)3)  is  formed  by  oxidizing  ferrous 
sulphate  with  nitric  acid.     When  ferric  sulphate  solution  is  mixed 


IRON  — NICKEL   AND   COBALT  349 

with  the  proper  quantity  of  potassium  (or  ammonium)  sulphate, 
iron  alum  (K2SO4.Fe2(SO4)3.24H2O)  is  formed.  It  is  a  pale  violet, 
crystallized  solid,  which  has  properties  like  ordinary  alum.  Iron 
alum  is  used  chiefly  as  a  mordant  (437,  438). 

469.  Iron  Sulphides.  —  Commercial  ferrous  sulphide  (FeS)  is  a 
black,  brittle,  metallic-looking  solid,  but  the  native  compound  is 
yellow  and  crystalline.  It  is  also  obtained  as  a  black  precipitate  by 
the  interaction  of  a  dissolved  ferrous  salt  and  ammonium  sulphide. 
It  is  made  on  a  large  scale  by  fusing  a  mixture  of  iron  and  sulphur, 
and  is  used  chiefly  in  preparing  hydrogen  sulphide.  Ferric  sulphide 
(iron  disulphide,  iron  pyrites,  pyrite,  FeS2)  is  one  of  the  commonest 
minerals.  It  is  a  lustrous,  metallic,  brass-yellow  solid.  Crystals  of 
pyrites  found  in  many  rocks  are  often  mistaken  for  gold  —  hence  the 
popular  name  "fool's  gold."  It  is  valueless  as  an  iron  ore,  but  large 
quantities  are  used  as  a  source  of  sulphur  in  making  sulphuric  acid 
(284). 

460.  Iron    Chlorides.  —  When   iron   interacts   with   hydrochloric 
acid,  ferrous  chloride  (FeCl2)  is  formed  in  solution.     Heated  in  the 
air  or  with  potassium  chlorate  or  nitric  acid,  it  is  changed  into  ferric 
chloride,  thus :  — 

2FeCl2  +       2HC1       +      O  2FeCl3  +  H2O 

Ferrous  Hydrochloric  Oxygen  Ferric  Water 

Chloride  Acid  Chloride 

Ferric  chloride  is  a  brownish  deliquescent  solid.  It  is  prepared  by 
passing  chlorine  into  ferrous  chloride  solution,  or  by  the  interaction 
of  iron  and  aqua  regia.  When  treated  with  nascent  hydrogen  or  an- 
other reducing  agent,  ferric  chloride  is  changed  into  ferrous  chloride. 
The  reaction  of  a  solution  of  ferric  chloride  is  acid,  owing  to  hydrolysis 
(165). 

461.  Ferrous    Carbonate  (FeCO3)  occurs  native  as  the  iron  ore 
siderite  (clay  iron  stone  or  spathic  iron  x>re).     The  typical  variety 
is  light  yellow  or  brown,  lustrous,  crystalline,  and  not  very  hard; 
but  many  kinds  are  impure,  and  the  properties  vary.     It  is  slightly 
soluble  in  water  containing  carbon  dioxide,  and  is  therefore  found  in 
some  mineral  springs. 

462.  Iron  Cyanides.  —  Iron  and  cyanogen  (CN)2,  with 
or   without   potassium,    form   several   compounds.     The 
most  important  is  potassium  ferrocyanide  (K4Fe(CN)e). 


350  CHEMISTRY 

It  is  a  lemon-yellow,  crystallized  solid,  containing  three 
molecules  of  water  of  crystallization.  Unlike  the  simple 
cyanogen  compounds  (e.g.  HCN  and  KCN),  it  is  not 
poisonous.  It  is  manufactured  by  fusing  iron  filings  with 
potassium  carbonate  and  nitrogenous  animal  matter 
(such  as  horn,  hair,  blood,  feathers,  and  leather).  The 
mass  is  extracted  with  water,  and  the  salt  is  separated 
by  crystallization.  Large  quantities  are  used  in  dyeing 
and  calico  printing,  and  in  making  bluing  and  potas- 
sium cyanogen  compounds.  Potassium  ferricyanide 
(K3Fe(CN)6)  is  a  dark  red  crystallized  solid,  containing 
no  water  of  crystallization.  It  is  often  called  red  prus- 
siate  of  potash.  It  is  manufactured  by  oxidizing  potas- 
sium ferrocyanide.  In  alkaline  solution  it  is  a  vigorous 
oxidizing  agent,  and  finds  extensive  use  in  dyeing.  It  is 
also  used  in  making  blue  print  paper. 

Blue  print  paper  is  made  by  coating  paper  with  a  mixture  of  solu- 
tions of  potassium  ferricyanide  and  ammonium  ferric  citrate  and  dry- 
ing in  a  dark  place.  In  the  sunlight  the  ferric  salt  is  partly  reduced 
and  forms  a  bronze  colored  deposit  by  interaction  with  the  potas- 
sium ferricyanide.  If  such  prepared  paper  is  covered  with  a  photo- 
graphic negative,  or  with  transparent  doth  on  which  lines  are  drawn 
in  black  ink,  and  placed  in  the  sunlight,  the  paper  is  acted  upon  only 
in  the  exposed  places.  Upon  washing,  the  exposed  parts  become 
blue,  and  the  covered  parts  white.  (See  Part  II,  Exp.  221.) 

463.  Tests  for  Iron.  —  Ferrous  salts  and  potassium 
ferricyanide  interact  in  solution  and  precipitate  ferrous 
ferricyanide  (Fe3(Fe(CN)6)2).  This  is  a  blue  solid  and 
is  often  called  Turnbull's  blue.  Ferric  salts  interact 
with  potassium  ferrocyanide  and  produce  ferric  ferro- 
cyanide (Fe4(Fe(CN)6)3).  This  precipitate  is  likewise  a 
dark  blue  solid,  and  is  called  Prussian  blue.  Prussian 
blue  is  extensively  used  in  dyeing  and  calico  printing, 


IRON  — NICKEL   AND   COBALT  351 

and  in  making  bluing.  By  the  above  reactions  ferrous 
and  ferric  salts  can  be  distinguished.  Besides  the  above 
tests  for  iron,  potassium  sulphocyanate  (KCNS)  produces 
a  red  solution  of  ferric  sulphocyanate  (Fe(CNS)3)  with 
ferric  salts,  but  leaves  ferrous  salts  unchanged.  (See 
Part  II,  Exps.  217,  218,  220.) 

Nickel  and  Cobalt 

464.  Nickel.  —  Small  amounts  of  metallic  nickel  are  found  in  mete- 
orites.    The  chief  ores  are  nickel-bearing  iron  sulphides,  which  are 
abundant  in  the  Sudbury  district,  Canada,  and  the  silicates  found  in 
New  Caledonia.     Nickel  is  a  silver-white  metal,  which  takes  a  bril- 
liant polish.     It  is  ductile,  hard,  malleable,  tenacious,  and  does  not 
tarnish  in  the  air.     Like  iron  it  is  attracted  by  a  magnet. 

Nickel  is  an  important  ingredient  of  coins  and  alloys  (399) .  The 
per  cent  of  nickel  is  12  in  the  United  States  cent  and  25  in  the  five- 
cent  piece.  German  silver  contains  from  20  to  25  per  cent  of  nickel, 
the  rest  being  copper  and  zinc.  Large  quantities  of  nickel  are  used  to 
coat  or  plate  other  metals,  especially  iron  and  brass.  Nickel  plating 
is  done  by  electrolysis,  as  in  silver  and  gold  plating,  though  the  electro- 
lytic solution  contains  ammonium  nickel  sulphate  (  (NH^aN^SO^), 
not  a  cyanide  (see  407,  413).  The  deposit  of  nickel  is  hard,  brilliant, 
and  durable.  Nickel  becomes  quite  malleable  if  a  little  magnesium 
is  added  to  the  molten  metal;  and  sheets  of  iron  upon  which  nickel  is 
welded  are  made  into  kitchen  utensils.  An  important  use  of  nickel 
is  in  the  manufacture  of  nickel  steel,  which  is  used  for  the  armor 
plate  and  turrets  of  battleships  (451). 

Nickel  forms  two  series  of  compounds  —  the  nickelous  and  the 
nickelic.  The  nickelous  are  more  common  and  many  of  them  are 
green.  The  test  for  nickel  is  the  formation  of  apple-green  nickelous 
hydroxide  (Ni(OH)j)  by  the  interaction  of  an  alkali  and  a  dissolved 
nickel  salt. 

465.  Cobalt,    Co,    generally   occurs   combined   with   arsenic   and 
sulphur  in  complex  minerals,  and  is  often  associated  with  nickel 
compounds.     It  is  a  lustrous  metal  with  a  reddish  tinge,  harder  than 
iron,  but  less  magnetic. 

Cobalt  forms  two  series  of  compounds  —  the  cobaltous  and  the  co- 


352  CHEMISTRY 

baltic.  The  cobaltous  compounds  are  more  common.  Cobaltous 
nitrate  (Co(NO3)2)  is  a  red  solid,  which  crystallizes  with  six  mole- 
cules of  water  of  crystallization.  The  hydrated  compound  loses  water 
readily  and  turns  blue  when  heated.  Some  cobalt  compounds  are 
used  to  color  glass,  porcelain,  and  paper,  especially  a  complex  com- 
pound known  as  smalt,  or  smalt  blue.  The  blue  color  produced  by 
fusing  cobalt  compounds  into  a  borax  bead  is  a  test  for  cobalt. 
Another  test  is  the  precipitation  of  yellow  potassium  cobaltinitrite 
(K3Co(NO2)6)  by  the  addition  of  potassium  chloride,  potassium 
nitrite,  and  acetic  acid  to  a  solution  of  a  cobaltous  compound. 

EXERCISES 

1.  Discuss  the  occurrence  of  iron.     Name  the  chief  ores.     Name  other 
native  compounds  of  iron.     What  proportion  of  the  earth's  crust  is  iron? 
Compare  with  the  abundance  of  other  elements. 

2.  Describe  a  blast  furnace.     Summarize  the  metallurgy  of  cast  iron. 

3.  Describe  the  manufacture  of  cast  iron.     State  its  general  composition, 
properties,  and  uses. 

4.  Apply  Exercise  3  to  wrought  iron. 

5.  State  the  properties  of  steel.     Compare  briefly  with  cast  and  wrought 
iron. 

6.  Describe  the  manufacture  of  steel  by  (a)  the  Bessemer  process  and 
(b)  the  open-hearth  process. 

7.  Essay  topics:    (a)  Production  and  transportation  of    iron    ore.     (b) 
Uses    of    steel,     (c)  Uses    of    nickel,     (d)  Meteorites,     (e)  Primitive    iron 
smelting.     (/)  Armor  plate,     (g)  Special  steels. 

8.  What  is  copperas,  rouge,   crocus,  iron    alum,  iron    pyrites,  green 
vitriol? 

9.  How  are  ferrous  changed  into  ferric  compounds,  and  vice  versa  ? 

10.  Practical  topics:    (a)  Cite  proofs  that  iron  is  widely  distributed. 
(b)  How  would  you  test  coal  ashes  for  iron? 

11.  Sketch  from  memory  a  vertical  section  of  (a)  a  blast  furnace,  (6)  a 
converter  in  operation,  (c)  an  open-hearth  furnace. 

12.  Starting  with  iron  how  would  you  prepare  in  succession  ferrous 
chloride,  FeCl3,  ferric  hydroxide,  ferric  chloride,  Fe4(Fe(CN)6)3? 

PROBLEMS 

1.  Calculate  the  per  cent  of  iron  in  (a)  olivine,  Mg2SiO4.Fe2SiO4,  (b) 
chalcopyrite,  CuFeS2,  (c)  potassium  ferrocyanide,  K4Fe(CN)6,  (d)  iron 
tetracarbonyl,  Fe(CO)4. 


IRON  — NICKEL  AND   COBALT  353 

2.  Calculate  the  weight  of  iron  in  (a)  70  tons  (2000  Ib.  each)  of  cop- 
peras, (b)  3  metric  tons  of  hematite  (95  per  cent  pure),  (c)  2  kg.  of  pyrite, 
(d)  1000  Ib.  of  magnetite,  (e)  1 75  milligrams  of  siderite. 

3.  Write  the  formulas  of  (a)  the  ferrous  and  (b)  the  ferric  salts  of  the 
following  acids:  Hydrobromic,  hydriodic,  carbonic,  nitric,  orthophosphoric, 
hydrosulphuric,  acetic. 

4.  Calculate    the   percentage  composition  of  (a)  the  three  oxides  and 
(b)  the  two  chlorides  of  iron,  and  show  how  the  two  sets  of  compounds 
illustrate  the  law  of  multiple  proportions. 

5.  Calculate  the  following:   (a)  the  weight  of  ferrous  carbonate  needed 
to  produce  25  1.  of  CO2  (standard  conditions);   (b)  the  weight  of  iron  formed 
by  the  interaction  of  hydrogen  and  220  gm.  of  Fe3O4;   (c)  the  weight  of  pure 
iron  that  can  be  made  from  1000  tons  of  iron  ore  (94  per  cent  of  hematite). 

6.  Complete  and  balance  the  following:    (a)  FeCl3  +  (NH4)2S  =  FeS 

+ ;      (b)  FeCl3  +  —  =  Fe(OH)3  +  NH4C1;      (c)  Fe  +  O  =  Fe2O3; 

(d)  K3Fe(CN)6  H = +  K2S04;      (e)  K4Fe(CN)6  + = 

+  KC1. 

7.  Write  the  formulas  of  all  the  hydroxides  of  iron,  cobalt,  and  nickel. 


CHAPTER  XXIX 

MAGNESIUM  —  ZINC  —  CADMIUM  —  MERCURY 

These  elements  form  a  natural  family,  though  the 
members  are  not  so  closely  related  as  in  the  alkali  and 
alkaline  earth  families. 

Magnesium 

466.  Occurrence.  —  Magnesium  is  never  found    free. 
In  combination  it  is  widely  distributed  and  very  abun- 
dant, constituting  about  2.5  per  cent  of  the  earth's  crust. 
Dolomite  is  magnesium  calcium  carbonate  (CaMg(CO3)2); 
it  forms  whole  mountain  ranges  in  the  Tyrol  and  vast 
deposits  in  many  regions.     Dolomite   closely  resembles 
marble   and   limestone   in   its   properties,  and   is    some- 
times called  magnesium  limestone.    Magnesium  carbonate 
(MgCO3)  is  also  abundant.     Many  of  the  Stassfurt  salts 
(377)   are  compounds  of  magnesium,  e.g.  kainite  (KC1. 
MgSO4.3H2O),  carnallite  (KCl.MgCl2.6H2O),  and  kieser- 
ite  (MgSO4.H2O).     Magnesium  is  also  a  constituent  of 
many  minerals  and  rocks,  such  as  serpentine,  talc,  soap- 
stone,  asbestos,  and  meerschaum.     The  sulphate  (MgSO4) 
and  chloride  (MgCl2)  are  found  in  sea  water  and  in  min- 
eral springs. 

467.  Preparation.  —  Magnesium   is  prepared   by  the   electrolysis 
of  fused  carnallite  (Fig.  93).     Carnallite  is  put  in  the  iron  vessel  C, 
which  is  the  cathode.     This  is  closed  by  the  air-tight  cover  through 
which  pass  the  pipes  D  D'  for  converting  an  inert  gas  into  and  out 
of  the  apparatus.     The  carbon  anode  A  dips  into  the  carnallite  and  is 


MAGNESIUM  —  ZINC  —  CADMIUM  —  MERCURY     355 


I 


E 


inclosed  by  the  porcelain  cylinder  B,  which  is  provided  with  a  pipe 
E  for  the  escape  of  the  chlorine  liberated  at  the  anode.  The  car- 
nallite  is  kept  fused  by  external  heat. 
When  the  current  passes,  the  chlorine 
escapes  through  E,  while  the  magne- 
sium rises  to  the  surface  of  the  fused 
carnallite  and  is  prevented  from  oxi- 
dizing by  the  inert  gas  supplied 
through  D.  The  porcelain  cylinder  B 
prevents  the  chlorine  from  escaping 
into  the  larger  vessel.  The  molten 
magnesium  is  carefully  removed  at 
intervals. 


Fig.  93.  —  Apparatus  for  the 
Manufacture  of  Magnesium 
by  the  Electrolysis  of  Car- 
nallite. 


468.  Properties. —  Magne- 
sium is  a  lustrous,  silvery  white 

metal.  It  is  a  light  metal,  the  specific  gravity  being 
about  1.75.  It  is  tenacious  and  ductile,  and  when  hot 
can  be  drawn  into  wire  or  pressed  into  ribbon,  the 
latter  being  a  common  commercial  form.  It  is  easily 
kindled  by  a  match,  melts  at  651°  C.,  and  can  be  cast. 
At  a  higher  temperature  it  volatilizes.  Heated  in  air, 
it  burns  with  a  dazzling  light,  producing  dense  white 
clouds  of  magnesium  oxide  (MgO)  together  with  a  little 
magnesium  nitride  (Mg3N2).  It  does  not  tarnish  in  dry 
air,  but  in  moist  air  it  is  soon  covered  with  a  film  of 
oxide.  It  liberates  hydrogen  from  acids.  When  heated 
in  dry  nitrogen,  it  forms  magnesium  nitride;  this  property 
was  utilized  by  Ramsay  in  the  separation  of  nitrogen 
from  argon  (129).  (See  Part  II,  Exp.  227.) 

469.  Uses.  —  The  light  from  burning  magnesium  affects  a  photo- 
graphic plate,  and  magnesium  powder  (alone  or  mixed  with  potassium 
chlorate)  is  used  in  taking  flashlight  photographs.     It  is  also  used 
for  signal  lights  and  fireworks.     Magnalium,  the  alloy  of  magnesium 
and  aluminium,  has  been  described  (399,  433). 


356  CHEMISTRY 

470.  Magnesium  Oxide  and  Hydroxide.  —  Magnesium 
oxide,  MgO,  is  a  white,  bulky  powder.     It  is  formed  when 
magnesium  burns  in  the  air,  but  it  is  manufactured  by 
gently   heating   magnesium    carbonate,    just   as   lime   is 
made  from  limestone.     It  is  often  called  magnesia,  or 
calcined  magnesia.    It  combines  slowly  with  water,  form- 
ing very  slightly  soluble  magnesium  hydroxide  (Mg(OH)2). 
Like  lime,  magnesia  withstands  a  very  high  temperature, 
and    is    therefore   used  in  making  fire  brick,   crucibles, 
furnace  linings,  and  for  other  purposes  requiring  a  refrac- 
tory  substance,    e.g.   as   an   ingredient   of   a    protective 
mixture  for  steam  pipes.      Both  magnesium  oxide  and 
hydroxide  are  used  as  a  medicine  for  dyspepsia  and  an 
antidote  for  poisoning  by  mineral  acids. 

471.  Magnesium   Sulphate,  MgSO4,  is  a  white  solid.     There  are 
several  crystallized  varieties.     The  native  salt  kieserite  (MgSO4.H2O) 
when  added  to  water  changes  into    Epsom  salts   (MgSO4.7H2O), 
which  is  the  commercial  form  of  magnesium  sulphate.     It  is  very 
soluble  in  water,  and  its  solution  has  a  bitter  taste.     Magnesium 
sulphate  makes  water  permanently  hard  (423).     It  is  used  as  a  med- 
icine, in  manufacturing  paints,  soap,  and  sulphates  of  sodium  and 
potassium,  as  a  fertilizer  in  place  of  gypsum,  and  as  a  coating  for 
cotton  cloth. 

472.  Magnesium  Chloride,   MgCl2,  is  a    white    solid. 
It   is    a   by-product   in    the   manufacture    of   potassium 
chloride    from    carnallite   (379).       The    crystallized    salt 
(MgCl2.6H2O)  is  very  deliquescent.     Magnesium  chloride 
undergoes  hydrolysis  with  hot  water,  forming  magnesium 
hydroxide  and  hydrochloric   acid.     If   water   containing 
magnesium   chloride  is   used  in   a   boiler,   the  insoluble 
magnesium  hydroxide  forms  a  hard  scale  on  the  boiler 
and  the  liberated  hydrochloric  acid  corrodes  the  metal. 
Hence  a  hard  water  containing  magnesium  chloride  (or 
sulphate)  should  be  softened  before  use  (423). 


MAGNESIUM  —  ZINC  —  CADMIUM  —  MERCURY     357 

If  disodium  phosphate  solution  and  ammonium  hydroxide  are 
added  to  a  solution  of  a  magnesium  compound,  the  precipitation  of 
ammonium  magnesium  phosphate  (NH4MgPO4)  serves  as  a  test  for 
magnesium. 

473.  Magnesium  Carbonate,  MgCO3,  occurs  native  as  magnesite, 
and  combined  with  calcium  carbonate  as  dolomite.     Like  the  corre- 
sponding calcium  compound,  it  forms  the  soluble  acid  carbonate  in 
water  containing  carbon  dioxide  (188).     The  commercial  salt  known 
as  magnesia  alba,  or  simply  magnesia,  is  a  complex  compound  (ordi- 
narily Mg  (OH)2.3MgCO3.3H2O).     Many  face  powders  consist  chiefly 
of  magnesia  alba.     Fluid  magnesia,   prepared  by  dissolving  mag- 
nesium carbonate  in  'water  containing  carbon  dioxide,  is  used  as  a 
medicine. 

Zinc 

474.  Occurrence.  —  Free   zinc   is   never   found.     The 
chief    ores    are    zinc    sulphide    (sphalerite,    zinc   blende, 
ZnS),  zinc  carbonate  (smithsonite,  ZnCO3),  zinc  silicates 
(hemimorphite,  Zn2SiO4.H2O,  willemite,  Zn2SiO4),  red  zinc 
oxide  (zincite,  ZnO),  and  franklinite  (Zn(Fe,  Mn)(FeO2)2). 
Zinc  ores  are  found  in  Missouri,  Kansas,  and  New  Jersey. 

475.  Metallurgy.  --  The  ores  are  roasted  and  then  re- 
duced by  heating  with  carbon  (Fig.  94).     The  reduction 


Fig.  94.  —  Retorts  for  Reduction  of  Zinc  Oxide  (Open  —  Right, 
Closed  — Left). 

is  conducted  in  earthenware  retorts  (A)  connected  with 
double  receivers;  at  first  the  zinc  condenses  in  C  as  a 
powder  known  as  zinc  dust,  somewhat  as  sulphur  forms 
flowers  of  sulphur,  but  when  this  receiver  becomes  hot, 
the  zinc  condenses  to  a  liquid  in  B,  from  which  it  is  drawn 


358  CHEMISTRY 

off  at  intervals  and  cast  into  bars  or  plates.  The  impure 
zinc  thus  obtained  is  called  spelter.  .It  is  freed  from 
carbon,  lead,  iron,  cadmium,  and  arsenic  by  repeated  dis- 
tillation; very  pure  zinc  is  obtained  by  the  electrolysis 
of  a  pure  zinc  salt. 

476.  Properties.  —  Zinc   is    a    bluish    white,    lustrous 
metal.     At  ordinary  temperatures  it  is  rather  brittle,  but 
at  ioo°-i5o°  C.  it  is  soft  and  can  be  rolled  into  sheets 
and  drawn  into  wire,  while  its  specific  gravity  rises  from 
6.9  to  7.2;  zinc  which  has  been  rolled  or  drawn  does  not 
become  brittle  upon  cooling.     Above  150°  C.  it  again  be- 
comes brittle.     It  melts  at  419.4°  C.  and  boils  at  918°  C. 
If  melted  zinc    is   poured   into  water,  it    forms    brittle 
lumps  called  granulated  zinc.     Heated  in  the  air  above 
its  melting  point,  zinc  burns  with  a  bluish  green  flame, 
forming  white  zinc  oxide  (ZnO).     Zinc  does  not  tarnish 
in  dry  air,  but  ordinarily  it  becomes  coated  with  a  dark 
film  which  is  essentially  a  basic  carbonate.     Commercial 
zinc  interacts   readily  with  acids   and  usually  liberates 
hydrogen;  pure  zinc  acts  very  slowly.     Like  aluminium, 
it  interacts  with  hot  solutions  of  sodium  and  potassium 
hydroxides  (431) ;  it  forms  zincates  and  hydrogen,  thus :  — 

2KOH          +  Zn  =      H2      +        K2ZnC2 

Potassium  Hydroxide        Zinc         Hydrogen         Potassium  Zincate 

Zinc  displaces  certain  metals  from  their  solutions  (414). 
(See  Part  II,  Exps.  184,  227.) 

477.  Uses.  —  Zinc    in    stick   or   plates   is   extensively 
used  as  an  electrode  in  electric  batteries.     Sheet  zinc  is 
used  as  a  lining  for  tanks.     The  chief  use  of  zinc  is  in 
making  galvanized  iron.     This  is  iron  coated  with  zinc 
and  is  made  by  dipping  clean  iron  into  melted  zinc.     The 
zinc  protects   the  iron  from  air  and  moisture.     Hence 


MAGNESIUM  —  ZINC  —  CADMIUM  —  MERCURY     359 

galvanized  iron  does  not  rust  easily  and  is  extensively 
used  for  netting,  wire,  roofs,  pipes,  cornices,  and  water 
tanks.  Zinc  dust  is  used  in  the  cyanide  process  of  ex- 
tracting gold  (411) .  Zinc  is  an  ingredient  of  many  alloys, 
e.g.  brass,  bronze,  and  German  silver  (399). 

478.  Compounds.  —  Native  zinc  oxide  is  red,  owing 
to  the  presence  of  manganese.  The  pure  oxide  (ZnO)  is 
white  when  cold  and  yellow  when  hot.  It  is  formed  when 
zinc  burns,  and  is  manufactured  in  this  way  or  by  heat- 
ing zinc  carbonate.  It  is  often  called  "zinc  white"  or 
"  Chinese  white,"  and  large  quantities  are  used  in  the 
manufacture  of  rubber  goods  and  to  make  a  white  paint 
which  is  not  discolored  by  sulphur  compounds  in  the 
atmosphere  (502).  Native  zinc  sulphide  is  yellow, 
brown,  or  black  on  account  of  impurities,  but  the  pure 
sulphide  (ZnS)  is  white.  The  latter  is  formed  as  a  jelly- 
like  precipitate  when  hydrogen  sulphide  is  passed  into  an 
alkaline  or  very  weak  acid  solution  of  a  zinc  salt.  Zinc 
sulphate  (ZnSO4)  is  formed  by  the  interaction  of  zinc  and 
dilute  sulphuric  acid.  Large  quantities  are  also  made  by 
roasting  the  sulphide  in  a  limited  supply  of  oxygen  and 
extracting  the  sulphate  with  water.  Thus  prepared,  it 
is  a  white,  crystallized  solid  (ZnSO4.7H2O)  which  efflo- 
resces in  the  air,  and  when  heated  to  100°  C.  loses  most 
of  its  water  of  crystallization.  The  crystallized  salt  is 
called  white  vitriol.  It  is  used  in  dyeing  and  calico  print- 
ing, as  a  disinfectant,  and  as  a  medicine.  It  is  poisonous. 
Zinc  chloride  (ZnCl2)  is  a  white,  deliquescent  solid.  It 
is  used  in  surgery,  and  also  as  a  constituent  of  a  mix- 
ture for  filling  teeth;  large  quantities  are  used  to  preserve 
wood,  especially  railroad  ties.  Zinc  hydroxide  (Zn(OH)2) 
is  formed  as  a  dull  white  precipitate  by  the  interaction 
of  sodium  or  potassium  hydroxide  and  a  solution  of  a 


360  CHEMISTRY 

zinc  salt.     An  excess  of  the  alkaline  hydroxide  changes 
zinc  hydroxide  into  a  zincate  (431) . 

479.  Tests  for  Zinc.  —  The  formation  of  the  sulphide  or  hydroxide, 
as  above  described,  serves  as  a  test  for  zinc.     A  green  incrustation  is 
produced  when  zinc  compounds  are  heated  on  charcoal  and  then 
moistened  with   cobaltous  nitrate  solution.     (Compare  434.)    (See 
Part  II,  Exp.  225.) 

480.  Cadmium,  Cd,  occurs  in  zinc  ores,  and  is  extracted  from  im- 
pure zinc.     It  is  white,  lustrous,  and  rather  soft.     Its  specific  gravity 
is  8.6,  and  it  melts  at  320.9°  C.     Cadmium  is  a  constituent  of  certain 
fusible  alloys  (356).    Wood's  metal,  for  example,  contains  12  per  cent 
of  cadmium.     The  most  important  compound  is  cadmium  sulphide 
(CdS).     This  is  a  bright  yellow  solid,  formed  by  adding  hydrogen 
sulphide  to  a  solution  of  a  cadmium  compound.     It  is  used  as  an 
artist's  color.     Its  formation  serves  as  the  test  for  cadmium. 

Mercury 

481.  Occurrence.  —  Native    mercury   is    occasionally 
found  in  minute  globules.     The  most  abundant  ore  and 
the  chief  source  of  mercury  is  mercuric  sulphide  (cinnabar, 
HgS).     The  ore  is  extensively  mined  in  Spain,  Austria, 
and   Italy;    in   the   United   States   large   quantities   are 
obtained  in  California  and  Texas. 

Mercury  has  been  known  for  ages  as  quicksilver.  The  Latin 
name,  hydrargyrum,  which  gives  the  symbol  Hg,  means  literally 
"  water  silver." 

482.  Metallurgy.  —  Mercury   is    readily  prepared    by 
roasting  cinnabar  in  a  current  of  air  or  with  calcium  oxide 
and  condensing  the  vapor.     The  equations  for  the  reac- 
tions are:  - 

HgS          +     02      =     Hg     +          S02 

Mercuric  Sulphide        Oxygen        Mercury        Sulphur  Dioxide 

4HgS  +  4CaO  =  4Hg  +  sCaS  +  CaSO4 

Mercuric         Calcium         Mercury         Calcium         Calcium 
Sulphide  Oxide  Sulphide        Sulphate 


MAGNESIUM  —  ZINC  —  CADMIUM  —  MERCURY     361 

Crude  mercury  is  freed  from  soot,  dirt,  and  mechanical  impurities 
by  pressing  it  through  linen  or  chamois  leather,  but  it  must  be  dis- 
tilled or  agitated  with  nitric  acid  (or  ferric  chloride)  to  remove  the 
dissolved  metals,  such  as  lead  or  zinc.  Mercury  is  sent  into  commerce 
in  strong  iron  flasks  holding  75  pounds. 

483.  Properties.  —  Mercury  is  a  bright,  silvery  metal, 
and  is  the  only  one  that  is  liquid  at  ordinary  tempera- 
tures.    It   solidifies   at   —  38.7°  C.  and  boils    at    about 
357°  C.     It  is  a  heavy  metal,  the  specific  gravity  being 
about  13.59.     Mercury  is  a  good  conductor  of  electricity. 
Mercury  does  not  tarnish  in  the  air,  unless  sulphur  com- 
pounds are  present.     At  a  high  temperature,  it  combines 
slowly  with  oxygen  to  form  the  red  oxide  (HgO).    Hydro- 
chloric acid  and  cold  sulphuric  acid  do  not  affect  it;  hot 
concentrated  sulphuric  acid  oxidizes  it,  and  nitric  acid 
changes  it  into  nitrates.  .  It  is  displaced  from  solution  by 
most  other  metals  (414).      (See  Part  II,  Exp.  229.) 

484.  Amalgams  are  alloys  of  mercury.     Amalgamated  zinc  is  used 
in  certain  electric  batteries  to  prevent  unnecessary  loss  of  the  zinc. 
Tin  amalgam  is  sometimes  used  to  coat  mirrors.     Amalgams  of  certain 
metals  are  used  as  a  filling  for  teeth.     Silver  and  gold  form  amalgams 
readily,  and  considerable  mercury  is  used  in  extracting  these  precious 
metals  from  their  ores  (404,  411).     Care  should  be  taken,  while  han- 
dling mercury,  not  to  let  it  come  in  contact  with  jewelry. 

485.  Uses.  —  Mercury  is   used   in    making  thermom- 
eters, barometers,  and  some  kinds  of  air  pumps.    Its  exten- 
sive use  in  extracting  gold  and  silver  has  been  mentioned. 
Considerable  is  used  in  preparing  certain  medicines  and 
explosives    (e.g.   mercury    fulminate,    which  is    used    in 
percussion  caps  and  cartridges). 

The  use  of  mercury  in  thermometers  depends  not  only 
on  the  fact  that  it  is  a  bright  liquid  between  a  wide  range 
of  temperature,  but  also  on  the  uniform  change  of  volume 
that  accompanies  a  change  of  temperature.  The  curve 


362 


CHEMISTRY 


13.6 
18.6 
13.4 
13  3 

\ 

s 

\ 

13.2 
13.1 
13.0 
12.9 
12,8 
£12.7 

Q       ( 

- 

\ 

- 

X 

x 

- 

\ 

L 

^ 

)                 100              200               300              400 
Temperature 

showing  the  relation  of  volume  and  temperature  is  almost 
a  straight  line  (Fig.  95),  that  is,  the  expansion  of  mercury 
is  regular. 

486.  Compounds  of  Mercury.  —  Mercury,  like  copper  and  iron, 
forms  two  classes  of  compounds  —  the  mercurous  and  the  mercuric. 

Mercuric  oxide  (HgO)  is 
a  red  powder,  produced 
by  heating  mercury  in 
air  or  by  heating  a  mix- 
ture of  mercury  and 
mercuric  nitrate.  The 
historical  importance  of 
this  compound  has  al- 
ready been  emphasized 
(10).  Mercurous  chlo- 
ride (HgCl)  is  a  white, 
tasteless  powder,  insolu- 
ble in  water.  It  is  formed 
when  a  chloride  and  mer- 
curous nitrate  interact  — 
a  test  for  mercury  in  mercurous  compounds.  Under  the  name  of 
calomel  it  is  extensively  used  as  a  medicine.  Mercuric  chloride 
(HgCl2)  is  a  white,  crystalline  solid,  soluble  in  water  and  in  alcohol. 
It  is  prepared  by  heating  a  mixture  of  mercuric  sulphate  and 
sodium  chloride.  Mercuric  chloride  is  a  violent  poison.  The  best 
antidote  is  the  white  of  a  raw  egg.  The  albumin  forms  an  in- 
soluble mass  with  the  poison,  which  may  then  be  removed  mechan- 
ically from  the  stomach.  The  common  name  of  mercuric  chloride 
is  corrosive  sublimate.  It  has  powerful  antiseptic  properties,  and 
is  extensively  used  in  surgery  to  protect  wounds  from  the  harmful 
action  of  germs;  taxidermists  sometimes  use  it  to  preserve  skins,  and 
it  has  many  serviceable  applications  as  a  medicine  and  disinfectant. 
It  is  usually  used  as  a  dilute  solution  (i  part  to  1000  parts  of  water). 
Mercuric  chloride  when  treated  with  a  reducing  agent,  such  as  stan- 
nous  chloride,  forms  at  first  white  mercurous  chloride  and  finally  a 
dark  gray  precipitate  of  finely  divided  mercury  —  the  test  for  mer- 
cury in  mercuric  compounds  (see  493).  The  equations  for  these 
reactions  are:  — 


Fig.  95.  —  Curve  Showing  Regular  Change  in 
Volume  of  Mercury  as  the  Temperature  is 
Changed. 


MAGNESIUM  —  ZINC  —  CADMIUM  —  MERCURY     363 

2HgCl2         +         SnCl2  2HgCl        +        SnCl4 

Mercuric  Chloride  Stannous  Chloride        Mercurous  Chloride        Stannic  Chloride 

2HgCl  +  SnCl2  =  2Hg  +  SnCl4 

Native  mercuric  sulphide  or  cinnabar  (HgS)  is  a  red,  crystalline  solid. 
When  hydrogen  sulphide  is  passed  into  a  solution  of  a  mercuric  salt, 
mercuric  sulphide  is  formed  as  a  black  powder;  this  variety,  when 
heated,  changes  into  red  crystals.  Vermillion  is  artificial  mercuric 
sulphide  prepared  by  various  processes.  It  has  a  brilliant  red  color, 
and  is  used  to  make  a  red  paint.  Mercurous  nitrate  (HgNO3)  and 
mercuric  nitrate  (Hg(NO3)2)  are  prepared  by  treating  mercury  re- 
spectively with  cold  dilute  nitric  acid,  and  with  hot  concentrated 
nitric  acid.  They  are  white,  crystalline  solids.  (See  Part  II,  Exp. 
226.) 

EXERCISES 

1.  Describe  the  manufacture  of  magnesium. 

2.  Starting  with  magnesium,  how  would  you  prepare  in  succession 
MgO,  magnesium  hydroxide,  MgCl2,  magnesium  carbonate,  magnesium 
oxide? 

3.  Write  equations  for  (a)  interaction  of  magnesium  and  sulphuric  acid 
and  (b)  heating  magnesium  in  nitrogen. 

4.  Essay  topics:   (a)  Asbestos,     (b)  Magnesia  as  a  refractory  material. 
(c)  Stassfurt  salts  containing  magnesium,    (d)  Dolomite,    (e)  Joseph  Black's 
investigations  of  "magnesia  alba."    (/)  Properties  of  zinc  and  aluminium. 
(g)  Corrosive  sublimate  and  other  germicides. 

6.  Name  the  chief  ores  of  zinc.     Describe  the  metallurgy  of  zinc. 

6.  Summarize  (a)  the  physical  and  (b)  the  chemical  properties  of  zinc. 

7.  Starting  with  zinc  how  would  you  prepare  in  succession  zinc  oxide, 
ZnCl2,  zinc  hydroxide,  Na2ZnO2,  zinc  sulphide,  ZnCl2,  ZnCO3,  zinc  oxide,  Zn? 

8.  Essay    topics:     (a)  Zinc    paints,     (b)  Galvanized    iron,     (c)  Amal- 
gams, (d)  History  of  mercury,     (e)  Cinnabar. 

9.  What  are  the  tests  for  zinc? 

10.  Describe  the  metallurgy  and  purification  of  mercury. 

11.  Practical  topics:   (a)  Suggest  a  proof  of  the  volatility  of  mercury  at 
ordinary  temperatures,      (b)  What  is  the  significance  of  "quick"  in  the 
word  quicksilver?     (c)  Etymology  of  amalgam,     (d)  In  what  respect  does 
mercury  resemble  bromine?     (e)  Name  three  metals  which  will  float  on 
mercury.     (/)  Why  is  mercury  used  in  a  barometer? 

12.  Describe  (a)  mercurous  chloride  and  (b)  mercuric  chloride.    What 
is  the  commercial  name  of  each?    The  use? 

13.  State  the  tests  for  mercury. 


364  CHEMISTRY 

14.  What  is  (a)  magnesia,  (b)  Epsom  salts,  (c)  galvanized  iron,  (d)  Chi- 
nese white,  (e)  white  vitriol,  (/)  calomel,  (g)  corrosive  sublimate? 

PROBLEMS 

1.  Calculate    the  per  cent  of   the  metallic  element  in  (a)  magnesium 
oxide,  zinc  oxide,  mercuric  oxide,  and  cadmium  oxide;    (b)  Epsom  salts, 
sphalerite,  cinnabar,  and  smithsonite;  (c)  Mg2P2O7,  H2Zn2SiO6,  Hg(NO3)2, 
CdS04. 

2.  Calculate  the  percentage  composition  of   mercurous   and   mercuric 
iodides,  and  show  that  these  compounds  illustrate  the  law  of  multiple  pro- 
portions.    (Use  exact  atomic  weights.) 

3.  Write  the  ordinary  and  the  ionic  equations  for  (a)  mercuric  chloride 
and  hydrogen  sulphide  form  mercuric  sulphide  and  hydrochloric  acid,  (b) 
magnesium  chloride  and  sodium  hydroxide  form  magnesium  hydroxide  and 
sodium  chloride,   (c)  zinc  hydroxide  and  sodium  hydroxide  form  sodium 
zincate  and  water,  (d)  cadmium  nitrate  and  hydrogen  sulphide  form  cad- 
mium sulphide  and  nitric  acid. 

4.  Write    the   formulas  of   the  following  compounds  by  applying  the 
principle  of  valence  or  by  utilizing  analogous  formulas  in  this  chapter: 
Magnesium  bromide,  magnesium  nitrate,  magnesium  sulphide,  zinc  chro- 
mate,   zinc  carbonate,   zinc  acetate,   zinc  phosphate   (ortho),   mercurous 
fluoride,  mercuric  sulphate,  mercurous  oxide,  cadmium  hydroxide. 

6.  The  annual  production  of  quicksilver  in  the  United  States  is  about 
20,600  flasks  of  75  pounds  each.  If  this  amount  was  transformed  without 
loss  into  corrosive  sublimate,  how  many  metric  tons  would  be  produced? 

6.  What  (a)  weight  of  mercury,   (b)   weight  of  sulphur  dioxide,  and 
(c)  volume  of  sulphur  dioxide  (standard  conditions)  can  be  obtained  from 
a  metric  ton  of  cinnabar  (60  per  cent  pure)? 

7.  Calculate    the    atomic  weights  of    magnesium,   mercury,  zinc,   or 
cadmium  from  the  following:    (a)  16.0263  gm.  of  MgO  give  47.8015  gm.  of 
MgSO4;  (b)  16.03161  gm.  of  zinc  give  20.2608  gm.  of  ZnO;  (c)  118.3938  gm. 
of  HgO  give  109.6308  gm.  of  mercury;  (d)  177.1664  gm.  of  mercuric  sulphide 
give  152.745  gm.  of  mercury;  (e)  23.3275  gm.  of  CdBr2  give  32.2098  gm.  of 
silver  bromide.     (Use  exact  atomic  weights.) 


CHAPTER   XXX 
TIN  —  LEAD 

Tin 

487.  Occurrence.  —  Tin  dioxide  (cassiterite,  tin  stone, 
SnO2)  is  the  only  available  ore.     It  is  not  widely  dis- 
tributed, and  is  found  chiefly  in  England  (at  Cornwall), 
Germany  (in  Bohemia  and  Saxony),  Australia,  Tasmania, 
and  the  East  Indian  Islands,  especially  Banca  and  Billiton. 
None  is  mined  in  the  United  States. 

488.  History.  —  Tin  is  one  of  the  oldest  metals.     Many  ancient 
bronzes  contain  tin.     The  Latin  word  stannum  gives  us  the  symbol 
Sn  and  the  terms  stannous  and  stannic. 

489.  Metallurgy. —  If   the   tin  ore   contains  sulphur   or   arsenic, 
these  impurities  are  removed  by  roasting.     The  tin  oxide  is  then 
reduced  by  heating  it  with  coal  in  a  reverberatory  furnace  (Fig.  90). 
The  simplest  equation  of  this  change  is  — 

Sn02     +     C     =  Sn  +       CO2 

Tin  Dioxide        Carbon        Tin        Carbon  Dioxide 

The  molten  tin  collects  at  the  bottom  of  the  furnace  and  is  drawn  off 
and  cast  into  bars  or  masses,  which  are  often  called  block  tin.  Usually 
it  is  purified  by  melting  it  slowly  on  a  hearth,  inclined  so  that  the  more 
easily  melted  tin  will  flow  down  the  hearth  and  leave  the  metallic 
impurities  behind.  This  tin  may  be  further  purified  by  stirring  the 
molten  metal  with  a  wooden  pole,  or  by  holding  billets  of  wood  be- 
neath its  surface.  The  impurities  which  are  oxidized  by  the  escaping 
gases  collect  as  a  scum  on  the  surface  and  are  removed. 

490.  Properties.  —  Tin    is    a    white,    lustrous    metal, 
which   does   not   tarnish   easily  in   air.     It  is   soft   and 
malleable,  and  can  be  readily  cut  and  hammered.     It  is 


366  CHEMISTRY 

softer  than  zinc,  but  harder  than  lead.  Its  specific 
gravity  is  about  7.3.  Tin  may  be  obtained  in  the  crys- 
talline form,  and  when  a  piece  of  such  tin  is  bent  it  makes 
a  crackling  sound,  which  is  probably  caused  by  the 
friction  of  the  crystals  upon  one  another.  It  melts  at 
231.9°  C.,  and  when  heated  to  a  higher  temperature  it 
burns,  forming  white  tin  oxide  (SnO2).  Ordinary  tin  if 
kept  below  about  20°  C.  changes  into  gray  tin,  which  is 
a  dull  looking  powder.  Sometimes  objects  containing 
tin,  such  as  organ  pipes,  medals,  and  statues,  disintegrate 
owing  to  the  formation  of  powdery  tin;  once  started, 
the  "tin  disease,"  as  it  is  called,  spreads  rapidly.  Con- 
centrated hydrochloric  acid  changes  it  into  stannous 
chloride  (SnCl2);  with  hot  concentrated  sulphuric  acid, 
it  forms  stannous  sulphate  (SnSO4)  and  sulphur  dioxide; 
and  concentrated  nitric  acid  oxidizes  it,  the  white,  solid 
product  being  known  as  metastannic  acid  (probably 
(H2SnO3)5).  Certain  metals  precipitate  tin  from  its  solu- 
tions often  as  a  grayish  black,  spongy  mass  filled  with 
bright  scales  (414).  (See  Part  II,  Exp.  234.) 

491.  Uses. --Tin  is  so  permanent  in  air,  weak  acids 
(like   vinegar   and   fruit   acids),  'and   alkalies   that  it  is 
extensively  used  as  a  protective  coating  for  metals.     The 
tin  plate  (sheet  tin,  or  simply  "tin")  is  made  by  dipping 
very  clean  sheet  iron  or  steel  into  molten   tin.     Thus 
coated,  it  is  made  into  tinware,  cans,  and  many  useful 
objects.     Copper  coated  with  tin  is  made  into  vessels  for 
cooking,  and  brass  coated  with  tin  is  made  into  pins. 
Tinned  iron  or  steel  does  not  rust  until  the  iron  is  exposed, 
and    then    the    rusting    proceeds    rapidly.     Tin    is    also 
hammered  into  thin  sheets  called  tin  foil,  though  much 
tin  foil  contains  lead. 

492.  Alloys.  —  Those  containing  a  minor  per  cent  of 


TIN  —  LEAD  367 

tin  are  bronze,  gun  metal,  type  metal,  and  fusible  alloys 
(399,  352,  356).  Speculum  metal  contains  about  30 
per  cent  of  tin.  Alloys  containing  considerable  tin  are 
Britannia  metal  (90  per  cent),  pewter  (75  per  cent),  and 
solder  (50  per  cent).  (Compare  499.) 

493.  Compounds  of  Tin.  —  Tin  forms  two  series  of  compounds 
—  the  stannous  and  stannic.     Stannic  oxide  or  tin  dioxide  (Sn02) 

has  already  been  mentioned  as  the  chief  ore  of  tin,  and  as  the  product 
formed  when  tin  is  burned.  The  artificial  oxide  is  faint  yellow  when 
hot  and  white  when  cold.  Stannous  chloride  (SnCl2)  is  formed  by 
the  interaction  of  hydrochloric  acid  and  tin.  From  the  concentrated 
solution  a  greenish  salt  crystallizes  (SnCl2.2H20),  known  as  tin  crystals 
or  tin  salt.  Stannous  chloride  can  be  readily  oxidized  to  stannic 
chloride  (SnCl4)  by  mercuric  chloride  solution.  The  equation  for 
this  change  is  — 

SnCl2       +       2HgCl2         =        SnCl4      +        2HgCl 

Stannous  Chloride        Mercuric  Chloride        Stannic  Chloride        Mercurous  Chloride 

By  an  extension  of  the  simplest  idea  of  oxidation  and  reduction  to 
include  the  negative  element  chlorine,  stannous  chloride  is  said  to 
be  oxidized  to  stannic  chloride  and  mercuric  chloride  to  be  reduced 
to  mercurous  chloride.  An  excess  of  stannous  chloride  reduces  the 
white  mercurous  chloride  to  gray  or  black  metallic  mercury;  this 
reaction  serves  as  a  test  for  tin.  (Compare  the  test  for  mercury  in 
mercuric  compounds,  486.)  Stannous  chloride  is  used  as  a  reducing 
agent  and  as  a  mordant  in  dyeing  and  calico  printing  (438).  Crys- 
tallized stannic  chloride  (SnCl4.sH2O),  known  commercially  as 
oxymuriate  of  tin,  is  also  used  as  a  mordant.  Tin  mordants  pro- 
duce brilliant  colors.  With  hydrogen  sulphide,  stannous  compounds 
form  brown  stannous  sulphide  (SnS),  while  stannic  compounds 
form  yellow  stannic  sulphide  (SnS2);  both  sulphides  dissolve  in 
ammonium  polysulphide,  owing  to  the  formation  of  soluble  sulpho- 
salts  of  tin.  (See  Part  II,  Exp.  232.) 

Lead 

494.  Occurrence.  —  The  most  abundant   ore   of   lead 
and  the  chief   commercial  source  of  the  metal  is  lead 


368  CHEMISTRY 

sulphide  (galena,  PbS).  Other  native  compounds  are  the 
carbonate  (cerussite,  PbCOs)  and  the  sulphate  (anglesite, 
PbSO4).  Lead  ore  is  found  in  the  United  States  in  the 
Middle  West  (Illinois,  Iowa,  Wisconsin,  and  Missouri), 
Colorado,  Idaho,  and  Utah.  Spain,  Mexico,  and  Ger- 
many produce  considerable. 

495.  History.  —  Lead  and  its  compounds  have  been  used  since  the 
dawn  of  history.  The  Chinese  have  used  it  for  ages  to  line  chests  in 
which  tea  is  stored  and  transported.  The  Romans,  who  obtained  it 
from  Spain,  called  it  plumbum  nigrum,  i.e.  black  lead,  and  used  it 
for  conveying  water  just  as  we  do  today.  The  symbol  Pb  comes  from 
plumbum. 

498.  Metallurgy.  —  Lead  is  obtained  from  galena  by 
several  processes,  (i)  Ores  rich  in  lead  are  roasted  in 
a  reverberatory  furnace  (Fig.  90)  until  a  part  of  the  sul- 
phide is  changed  into  lead  oxide  and  lead  sulphate.  The 
equations  for  these  changes  are  - 

2PbS       +   302    =     2PbO     +         2S02 

Lead  Sulphide        Oxygen        Lead  Oxide        Sulphur  Dioxide 


PbS        +    202    =       PbSO4 

Lead  Sulphide        Oxygen        Lead  Sulphate 

The  air  is  then  shut  off  and  the  mixture  of  the  three  lead 
compounds  is  heated  to  a  higher  temperature.  By  this 
operation  the  lead  sulphide  interacts  with  the  other  lead 
compounds,  forming  lead  and  sulphur  dioxide,  thus  - 

2PbS         +     PbS04     +  2PbO  =  5Pb  +        3S02 

Lead  Lead  Lead          Lead  Sulphur 

Sulphide  Sulphate  Oxide  Dioxide 

(2)  Ores  poor  in  lead  are  roasted  with  iron,  which  com- 
bines with  the  sulphur,  leaving  the  lead  free,  thus:- 
PbS         +  Fe  =  Pb  +       FeS 

Lead  Sulphide        Iron        Lead        Iron  Sulphide 


TIN  —  LEAD  369 

(3)  Ores  rich  in  silver  are  first  roasted  and  then  heated 
with  a  mixture  of  coke,  limestone,  and  iron  ore.  The 
lead,  gold,  silver,  and  other  metals  collect  as  a  liquid  in 
a  receptacle  in  the  lower  part  of  the  furnace. 

Lead  produced  by  these  processes  is  impure  and  must 
be  refined.  This  is  done  by  first  heating  the  metal  in  a 
reverberatory  furnace  (Fig.  90)  to  oxidize  most  of  the 
copper,  arsenic,  and  antimony,  and  then  treating  the 
alloy  of  lead,  gold,  silver,  etc.,  by  the  Parkes  process 
(404),  or  by  an  electrolytic  process  somewhat  like  that 
used  for  copper  (398).  In  refining  lead  by  electrolysis, 
the  cathode  is  a  sheet  of  pure  lead,  the  anode  is  a  heavy 
plate  of  impure  lead,  and  the  electrolytic  solution  is  a 
mixture  of  lead  fluosilicate  (PbSiF6)  and  gelatin.  When 
the  current  passes,  pure  lead  is  deposited  upon  the  cathode 
and  most  of  the  other  metals  remain  attached  to  the 
remnant  of  the  anode;  subsequently  the  gold  and  silver 
as  well  as  bismuth  are  recovered. 

497.  Properties.  —  Lead  is  a  bluish  metal.  When 
scraped  or  cut,  it  has  a  brilliant  luster,  which  soon  dis- 
appears, owing  to  the  formation  of  a  film  of  oxide.  It  is 
a  soft  metal,  and  can  be  scratched  with  the  finger  nail. 
It  discolors  the  hands,  and  when  drawn  across  a  rough 
surface  it  leaves  a  black  mark.  For  this  reason  it  is 
sometimes  called  black  lead  (172).  Lead  is  not  tough 
nor  very  ductile  and  malleable,  though  it  can  be  made 
into  wire,  rolled  into  sheets,  and  pressed  while  soft  into 
pipe.  It  is  a  heavy  metal,  its  specific  gravity  being 
11.34;  with  the  exception  of  mercury,  it  is  the  heaviest 
of  the  familiar  metals.  It  melts  at  327.4°  C.  Lead 
when  heated  strongly  in  air  changes  into  oxides  (mainly 
the  monoxide,  PbO).  Hydrochloric  and  sulphuric  acids 
have  little  effect  upon  compact  lead.  Nitric  acid  changes 


370  CHEMISTRY 

it  into  lead  nitrate  (Pb(N03)2).  Acetic  acid  (or  vinegar) 
and  acids  from  fruits  and  vegetables  change  it  into 
soluble,  poisonous  compounds;  hence  cheap  tin-plated 
vessels,  which  sometimes  contain  lead,  should  never  be 
used  in  cooking.  Certain  metals  precipitate  lead  from 
its  solutions  as  a  grayish  mass,  which  often  has  a  beautiful 
treelike  appearance  (414).  (See  Part  II,  Exp.  234.) 

498.  Uses.  —  Lead  is  extensively  used  as  pipe.     Lead 
pipe  is  not  only  used  to  convey  water  to  and  from  parts 
of  buildings,  but  as  a  sheath  for  copper  wires,  both  over- 
head and  underground.     Sheet    lead   is    used    to    cover 
roofs  and  to  line  sinks,  cisterns,  and  the  cells  employed 
in  many  electrolytic  processes.     The  lead  chambers  and 
some  evaporating  pans  used  in  manufacturing  sulphuric 
acid  are  made  of  sheet  lead.     Shot  and  bullets  are  lead 
(alloyed  with  a  little  arsenic).     Spongy  lead  is  used  in 
preparing  the  plates  of  storage  batteries. 

499.  Alloys.  —  Alloys  containing  considerable  lead  are 
type  metal,  solder,  Babbitt  metal  and  pewter  (399,  352, 
492).     Most  fusible  metals  contain  lead  (356). 

500.  Lead  Oxides.  —  There  are  three  important  oxides. 
Lead  monoxide  (PbO)  is  a  yellowish  powder  known  as 
massicot,  or  a  buff-colored  crystalline  mass  called  lith- 
arge.    It  is  formed  by  heating  melted  lead  in  a  current 
of  air.     It  is  made  this  way,  though  considerable  is  ob- 
tained as  a  by-product  in  separating  silver  from  lead. 
Large   quantities   are  used  in  preparing  some   oils  and 
varnishes,  flint  glass,  lead  compounds,  and  as  a  glaze  for 
pottery.      Lead   tetroxide  (red   lead  or  minium,  Pb3O4) 
is  a  red  powder,  varying  somewhat  in  color  and  composi- 
tion.    It  is  prepared  by  heating  lead  (or  lead  monoxide) 
to  about  350°  C.     It  is  used  in  making  flint  glass;  pure 
grades  are  made  into  artists'  paint,  but  the  cheap  variety 


TIN  —  LEAD  371 

is  used  to  paint  structural  iron  work  (bridges,  gasometers, 
etc.),  hulls  of  vessels,  and  agricultural  implements.  A 
mixture  of  linseed  oil  and  red  lead  is  used  in  plumbing 
and  gas  fitting  to  make  joints  tight.  Orange  mineral 
has  about  the  same  composition  as  red  lead,  though  its 
color  is  lighter;  its  uses  are  the  same.  Lead  dioxide 
(lead  peroxide,  Pb02)  is  a  brown  powder  formed  by 
treating  lead  tetroxide  with  nitric  acid  or  by  the  action 
of  chlorine  on  an  alkaline  solution  of  lead  acetate.  It  is 
used  in  storage  batteries. 

501.  Lead  Carbonate,  PbCO3,  is  found  native  as  the 
transparent,  crystallized  mineral  cerussite.  It  is  obtained 
as  a  white  powder  by  adding  sodium  bicarbonate  solution 
to  a  solution  of  a  lead  salt.  Sodium  and  potassium  car- 
bonates, however,  produce  basic  lead  carbonates.  The 
most  important  of  these  basic  carbonates  has  the  com- 
position corresponding  to  the  formula  2PbCO3.Pb(OH)2, 
and  is  known  as  white  lead.  It  is  a  heavy,  white  powder 
which  mixes  well  with  linseed  oil,  and  is  used  extensively 
as  a  white  paint  and  as  the  basis  of  many  colored  paints. 
White  lead  paint  covers  a  surface  well  and  dries  to  a  good 
finish.  But  it  darkens  on  exposure  to  hydrogen  sulphide, 
which  is  often  present  in  the  air  of  cities  (274) .  In  recent 
years  other  paint  bodies,  as  the  solids  are  called,  have 
been  mixed  with  or  substituted  for  white  lead,  e.g.  zinc 
oxide,  kaolin,  barium  sulphate,  and  lithophone.  These 
are  white  solids  which  do  not  darken  in  the  air,  and  they 
often  improve  the  paint  in  other  ways. 

602.  Manufacture  of  White  Lead.  —  White  lead  is  manufactured 
by  several  processes.  The  Dutch  process  is  the  oldest,  having  been 
used  as  early  as  1622.  It  is  essentially  the  same  today,  though  many 
details  have  been  improved.  Perforated  disks  of  lead  are  put  in  earth- 
enware pots  which  have  a  separate  compartment  at  the  bottom 


372 


CHEMISTRY 


containing  a  weak  solution  of  acetic  acid  (Fig.  96).  These  pots 
are  arranged  in  tiers  in  a  large  building,  and  spent  tan  bark  is  placed 
between  each  tier.  The  building  is  now  closed  except  openings  for 

the  entrance  and  exit  of  air  and 
steam.  The  heat  volatilizes  the 
acetic  acid  which  changes  the  lead 
into  lead  acetate.  The  tan  bark 
ferments  and  liberates  carbon  diox- 
ide, which  converts  the  lead  acetate 
into  basic  lead  carbonate  or  white 
lead.  The  operation  is  allowed  to 
proceed  until  the  lead  is  entirely 
transformed  —  sixty  to  one  hundred 
days.  Commercial  white  lead  is 
manufactured  by  other  processes, 
the  products  varying  somewhat  with 
the  process. 

503.  Other  Compounds  of   Lead. 


Fig.  96.  —  Earthenware  Vessel 
Containing  Lead  Disks  to  be 
Made  into  White  Lead.  Disk 
Before  (Lower)  and  After 
(Upper)  Corrosion. 


—  Native  lead  sulphide  (PbS)  is  the  mineral  galena,  the  chief  ore  of 
lead.  It  resembles  lead  in  appearance,  but  is  harder  and  is  usually 
crystallized  as  cubes,  octahedrons,  or  their  combinations  (Fig.  97). 
It  is  obtained  as  a  black  precipitate  by  the  interaction  of  hydrogen 
sulphide  (or  other  soluble  sulphides)  and  a  solution  of  a  lead  salt. 
Its  formation  is  the  test  for  lead.  Lead  chloride  (PbQ2)  is  a  white 


Fig.  97. —  Galena  Crystals  (Cube,  Octahedron  and  Cube,  Octahedron). 

solid  formed  by  adding  hydrochloric  acid  or  a  soluble  chloride  to  a 
cold  solution  of  a  lead  salt.  It  dissolves  in  hot  water.  Lead  sul- 
phate (PbSO4)  is  a  white  solid,  formed  by  adding  sulphuric  acid  or 
a  soluble  sulphate  to  a  solution  of  a  lead  salt.  It  is  very  slightly 
soluble  in  water,  but  soluble  in  concentrated  sulphuric  acid,  hence 
crude  sulphuric  acid  often  contains  lead  sulphate.  Lead  nitrate 
(Pb(NO3)2)  is  a  white  crystalline  solid  formed  by  dissolving  lead  (or 


TIN  —  LEAD  373 

lead  monoxide)  in  dilute  nitric  acid.  Lead  acetate  (Pb(C2H3O2)2) 
is  a  white,  crystalline  solid  formed  by  the  action  of  acetic  acid  upon 
lead  or  lead  oxide  (PbO).  Lead  Chromate  (PbCrO4)  is  a  yeUow  solid 
formed  by  adding  a  solution  of  a  lead  compound  to  a  solution  of  potas- 
sium chromate  or  potassium  dichromate.  It  is  sometimes  called 
chrome  yellow  and  is  used  as  a  pigment.  Its  formation  serves  as  a 
test  for  lead.  (See  Part  II,  Exp.  233.) 

504.  Cerium  (Ce)  and  Thorium  (Th)  are  members  of  a  family  in 
the  same  periodic  group  as  tin  and  lead.  They  are  constituents  of 
rare  minerals.  Their  compounds  are  prepared  from  monazite  sand. 
A  mixture  of  the  oxides  of  thorium  and  cerium  composes  the  Welsbach 
mantle  (213).  In  making  mantles,  a  cotton  bag  is  dipped  into  a 
solution  of  thorium  and  cerium  nitrates  and  then  burned.  The  cotton 
is  consumed  and  the  nitrates  are  changed  into  a  firm  mass  of  oxides. 
(See  Part  II,  Exp.  102.)  Thorium  is  a  radioactive  element  (526). 

EXERCISES 

1.  Describe  the  metallurgy  of  tin  and  of  lead. 

2.  Summarize  the  properties  of  tin  and  of  lead.     State  their  uses. 

3.  Describe  three  alloys  which  contain  large  proportions  of  tin.     Name 
several  alloys  containing  a  minor  proportion  of  tin. 

4.  Essay  topics:    (a)   History  of  tin.     (b)   Tin  disease,     (c)  Tin  plate 
industry,     (d)  Tin  mordants,     (e)  Tin  foil  and  its    substitutes.     (/)  Re- 
covery of  tin  from  tin  scrap. 

PROBLEMS 

1.  Calculate  the  weight  of  (a)  lead  in  250  gm.  of  PbO2  and  (b)  of  tin 
in  250  gm.  of  SnO2. 

2.  How  many  gm.  of  lead  in  (a)  200  gm.  of  galena,  (b)  i  kg.  of  litharge, 
(c)  a  metric  ton  of  red  lead,  (d)  750  gm.  of  cerussite? 

3.  Write  the  formulas  of  the  following  compounds  by  applying  the 
principle  of  valence :  Lead  fluoride,  lead  acetate,  lead  dichromate,  stannous 
iodide,  stannic  bromide,  stannous  sulphide,  stannic  sulphide. 

4.  Calculate  the  atomic  weight  of  lead  or  tin  from  the  following:    (a) 
16.2956  gm.  of  lead  give  17.554  gm.  of  PbO;  (b)  4.9975  gm.  of  PbCl2  require 
3.881  gm.  of  silver  to  precipitate  the  lead;    (c)  25  gm.  of  tin  give  31.8  gm. 
of  stannic  oxide  (the  specific  heat  of  tin  is  .055);   (d)  29.42  gm.  of  tin  unite 
with  35.4  gm.  of  chlorine,  and  the  vapor  density  of  the  compound  is  8.303. 

5.  Complete  and  balance  the  following:  (a)  SnCl2  +  HgCl2  =  SnCl4  + 

;    (b)  Sn  +  HNO3  =  H2SnO3  +  4NO  H ;    (c)  SnCl2  +  H2S  =  SnS 

+ ;    (d)  Pb(N03)2  +  HC1  =  PbCl2  + . 


CHAPTER  XXXI 
CHROMIUM  —  MANGANESE 

Chromium 

505.  Occurrence.  —  Metallic  chromium  is  never  found 
free.     Its  chief  ore  is  ferrous  chromite  (chrome  iron  ore, 
Fe(Cr02)2).     Native  lead  chromate  (crocoite,  PbCrO4)  is 
less  common.     Traces  of  chromium  occur  in  many  miner- 
als and  rocks,  e.g.  emerald  and   serpentine,   and  verde 
antique  marble.     Chromite  is  mined  chiefly  in   Greece, 
New  Caledonia,  New  South  Wales,  and  Turkey. 

506.  Preparation,  Properties,  and  Uses.  —  Chromium  is  produced 
in  large  quantities  by  reducing  chromic  oxide  with  granulated  alu- 
minium (see  Thermit,  432). 

Chromium  is  a  silvery,  crystalline,  hard,  and  brittle  metal.  Its 
specific  gravity  is  about  6.9  and  its  melting  point  is  1510°  C.  It 
is  not  oxidized  by  air  at  ordinary  temperatures. 

Chromium  is  used  to  harden  steel,  especially  the  kind  that  is  made 
into  armor  plate,  projectiles,  safes,  vaults,  and  parts  of  the  machines 
used  to  crush  gold-bearing  quartz.  This  hardened  steel  is  called 
chrome  steel.  It  forms  other  alloys  which  are  very  hard.  One 
commercial  form  of  chromium  is  an  alloy  of  65  to  80  per  cent  chro- 
mium, a  little  carbon,  and  the  rest  iron;  this  alloy  is  called  ferro- 
chrome  (451).  Another  alloy,  called  nichrome,  contains  nickel. 

507.  Potassium   Chromate  and  Potassium  Bichromate 
(or  Bichromate).  —  Potassium    chromate   (K2CrO4)  is  a 
lemon-yellow,   crystalline    solid,   very  soluble   in  water. 
Acids  change  it  into  the  dichromate,  thus :  - 

2K2CrO4  +  H2SO4  =  K2Cr207  +    K2SO4    +  H2O 

Potassium          Sulphuric        Potassium          Potassium         Water 
Chromate  Acid  Dichromate          Sulphate 


CHROMIUM  —  MANGANESE  375 

Potassium  Dichromate  (K2Cr207)  is  a  red  solid.  It 
forms  large  crystals  which  are  anhydrous.  It  is  less 
soluble  in  water  than  potassium  chromate.  Alkalies 
change  it  into  a  chromate,  thus  - 

K2Cr2O7    +   2KOH  =  2K2CrO4  +  H2O 

Potassium  Potassium         Potassium         Water 

Dichromate  Hydroxide        Chromate 

Potassium  dichromate  is  used  in  dyeing,  calico  printing, 
and  tanning,  in  bleaching  oils,  and  in  manufacturing 
other  chromium  compounds  and  dyestuffs.  Its  uses 
depend  mainly  upon  the  fact  that  it  is  an  oxidizing  agent. 
When  hydrochloric  acid  is  added  to  potassium  dichromate, 
the  oxygen  of  the  dichromate  oxidizes  the  hydrogen  of 
the  acid  and  liberates  chlorine,  thus:  - 

K2Cr2O7  +     I4.HC1     =  2KC1    +2CrC]3  +  3C12    +7H2O 

Potassium       Hydrochloric      Potassium      Chromic       Chlorine       Water 
Dichromate  Acid  Chloride        Chloride 

If  an  oxidizable  substance  is  present,   such  as  organic 

matter,   alcohol,   or   a  ferrous   compound,   it  is  quickly 

oxidized;  the  equation  for  the  reaction  in  the  case  of 
ferrous  sulphate  is  - 

K2Cr2O7  +  7H2SO4  +  6FeSO4  =  3Fe(SO4)3  +  Cr2(SO4)3  +  K2SO4  +  7H2O 

Ferrous  Ferric  Chromic 

Sulphate  Sulphate  Sulphate 

Potassium  chromate  and  dichromate  are  manufactured  from 
chrome  iron  ore.  The  crushed  ore  is  mixed  with  lime  and  potassium 
carbonate,  and  roasted  in  a  reverberatory  furnace;  air  is  freely  ad- 
mitted and  the  mass  is  frequently  raked.  By  this  operation  the  ore 
is  transformed  into  a  mixture  of  calcium  and  potassium  chromates. 
The  mass  is  cooled,  pulverized,  and  treated  with  a  hot  solution  of 
potassium  sulphate,  which  changes  the  calcium  chromate  into 
potassium  chromate.  The  potassium  chromate  is  changed  by  sul- 
phuric acid  into  potassium  dichromate;  the  latter  is  purified  by 
recrystallization  from  water. 

Potassium  chromate  is  also  formed  as  a  yellow  mass  by  fusing  on 


376  CHEMISTRY 

porcelain  or  platinum  a  mixture  of  a  chromium  compound,  potassium 
carbonate,  and  potassium  nitrate.  When  the  mass  is  boiled  with 
acetic  acid  to  decompose  the  carbonate  and  expel  carbon  dioxide, 
and  then  added  to  a  lead  salt  solution,  yellow  lead  chromate  is  formed. 
This  experiment  is  often  used  as  a  test  for  chromium.  (See  Part  II, 
Exps.  239,  240.) 


508.  Chrome  Alum,  K^Q^SO^^H^O,  is   a  purple, 
crystalline    solid.     It   is    analogous   in  composition  and 
similar  in  properties  to  ordinary  alum,  but  it  contains 
chromium  instead  of   aluminium  (437).     It  can  be  pre- 
pared by  mixing  potassium  and  chromium  sulphates  in 
the  proper  proportion,  or  by  passing  sulphur  dioxide  into 
a  solution  of  potassium  dichromate  containing  sulphuric 
acid.     Chrome  alum  is  used  as  a  mordant  in  dyeing  and 
calico  printing.     It  is  also  used  in  tanning  because  it 
acts  like  other  tanning  materials  in  much  less  time. 

509.  Lead  Chromate,  PbCr04,  is  a  bright  yellow  solid,  formed  by 
adding  potassium  chromate  or  dichromate  to  a  solution  of  a  lead 
salt.     Its  formation  is  a  test  for  chromium.     It  is  known  as  chrome 
yellow  and  is  used  in  making  yellow  paint  (504). 

510.  Other  Compounds  of  Chromium.  —  Three  other  chromium 
compounds  should  be  mentioned.     Chromic  oxide  (Cr2O3)  is  a  bright 
green  powder  prepared  by  heating  chromic  hydroxide  (Cr(OH)3),  and 
is  the  basis  of  the  chrome  green  pigments  used  to  color  glass  and  porce- 
lain.    When  chromium  compounds  are  heated  with  borax  they  color 
the  bead  green,  owing  to  the  formation  of  this  oxide.     There  are 
several  chromic  hydroxides.     The  typical  one  is  a  bluish  solid  formed 
by  the  interaction  of  a  chromic  compound  (e.g.  chrome  alum)  and  an 
alkaline  hydroxide,  carbonate,  or  sulphide.     The  chromic  hydroxide, 
which  is  always  precipitated,  is  soluble  in  an  excess  of  sodium  (or 
potassium)  hydroxide.     That  is,  it  is  changed  into  a  soluble  chro- 
mite,  just  as  aluminium  hydroxide  forms  soluble  aluminates.     Unlike 
aluminates,  however,  the  chromites  are  changed  back  into  chromic 
hydroxide  by  boiling.     When  concentrated  sulphuric  acid  is  added 
to   a  saturated  solution  of  potassium   dichromate   (or   chromate), 
chromium  trioxide   (CrO3)   separates  as  long,  bright  red  crystals; 


CHROMIUM  —  MANGANESE  377 

this  oxide  is  sometimes  called  chromic  acid.     It  is  a  vigorous  oxidizing 
agent. 

511.  Molybdenum  (Mo),  Tungsten  (W),  and  Uranium  (U)  are  rare 
metallic  elements  in  the  same  periodic  group  as  chromium.     Ammo- 
nium molybdate  (  (NH4)2MoO4)  is  used  to  detect  and  determine  phos- 
phorus in  fertilizers  (340).     Tungsten  is  used  to  harden  steel  and  as 
a  filament  of  electric  light  bulbs;  sodium  tungstate  is  used  for  mak- 
ing cloth  fireproof.     Uranium   compounds  are  obtained  chiefly  from 
pitchblende  and  uraninite.     Salts  of  uranium  (e.g.  sodium  uranate, 
Na2U2O7.6H20)  are  used  in  making  fluorescent  glass;    such  glass  is 
green  by  transmitted  light  and  yellow  by  reflected  light.     Uranium 
is  a  radioactive  element  (526) . 

Manganese 

512.  Occurrence.  —  This   metal   is  not  lound  free  in 
nature,   but  its  oxides   and  hydroxides  are  widely  dis- 
tributed and  rather  abundant.     The  chief  compound  is 
manganese  dioxide  (pyrolusite,  MnO2). 

513.  Preparation,  Properties,  and  Uses.  —  Manganese  is  prepared 
by  heating  manganese  dioxide  with  charcoal  in  an  electric  furnace, 
or  by  reducing  the  oxide  with  aluminium  powder  (432).     The  metal 
is  grayish,  hard,  and  brittle.     It  melts  at  1225°  C. 

514.  Alloys  of  manganese  and  iron  are  extensively  used  in  the  man- 
ufacture of  steel  (449,  451).     Spiegel  iron  contains  from  5  to  20  per 
cent  of  manganese,  while  ferromanganese  contains  20  per  cent  or 
more. 

515.  Manganese  Dioxide,  MnO2,  is  the  most  abundant 
and  important  compound.     It  is  a  black  solid  and  is  often 
called  black  oxide  of  manganese.     When  heated  it  yields 
oxygen;    and  when  heated  with  hydrochloric  acid  the 
two   compounds  interact,   forming  manganous   chloride, 
chlorine,  and  water,  thus  - 

MnO2    +      4HC1      =    MnCl2    +     C12     +  2H2O 

Manganese        Hydrochloric        Manganese        Chlorine         Water 
Dioxide  Acid  Chloride 


378  CHEMISTRY 

It  colors  glass  and  borax  a  beautiful  amethyst,  and  is 
often  used  in  glass  making  to  neutralize  the  green  color 
caused  by  impurities.  Large  quantities  are  used  in  the 
manufacture  of  oxygen,  chlorine,  glass,  and  manganese 
alloys  and  compounds. 

516.  Potassium  Permanganate,  KMn04,  is  a  dark 
purple,  glistening,  crystalline  solid,  though  the  crystals 
sometimes  appear  black  with  a  greenish  luster.  It  is 
very  soluble  in  water,  and  the  solution  is  red,  purple,  or 
black,  according  to  the  concentration.  Potassium  per- 
manganate gives  up  its  oxygen  readily  and  is  frequently 
used  as  an  oxidizing  agent.  It  is  also  used  as  a  disin- 
fectant, as  a  medicine,  in  bleaching  and  dyeing,  in  color- 
ing wood  brown,  and  in  purifying  gases,  such  as  hydrogen, 
ammonia,  and  carbon  dioxide.  (See  Part  II,  Exp.  244.) 

The  uses  of  potassium  permanganate  depend  mainly  upon  its  oxi- 
dizing power.  With  sulphuric  acid  the  action  is  represented  thus:  — 


2KMn04  +  3H2S04  =    50    +  2MnSO4  +  K2S04 

Potassium  Sulphuric        Oxygen        Manganese        Potassium        Water 

Permanganate  Acid  Sulphate  Sulphate 

The  liberated  oxygen  oxidizes  any  organic  matter  present,  and  the 
solution  becomes  brown  or  colorless,  owing  to  the  reduction  of  the 
potassium  permanganate  into  colorless  compounds. 

617.  Other  Compounds  of  Manganese.  —  Three  manganous  com- 
pounds are  important,  the  chloride  (MnCl2),  the  sulphate  (MnSO4), 
and  the  sulphide  (MnS).  The  chloride  and  sulphate  are  pink,  crys- 
talline salts.  The  sulphide  is  a  flesh-colored  precipitate  formed  by 
adding  ammonium  sulphide  to  the  solution  of  a  manganous  salt, 
the  color  distinguishing  it  from  all  other  sulphides;  its  formation 
serves  as  a  test  for  manganese.  Potassium  manganate  (K2MnO4) 
is  obtained  as  a  green  mass  by  fusing  a  mixture  of  a  manganese  com- 
pound, potassium  hydroxide  (or  carbonate),  and  potassium  nitrate. 
Its  formation  on  a  small  scale  constitutes  the  test  for  manganese. 
Sodium  manganate  (Na2MnO4)  is  used  in  solution  as  a  disinfectant. 
In  manganates  (and  also  in  permanganates)  manganese  acts  as  a 
non-metal.  (See  Part  II,  Exp.  241.) 


CHROMIUM—  MANGANESE  379 


EXERCISES 

1.  Describe  the  preparation  of   chromium  and  of   manganese   by  the 
aluminothermic  method. 

2.  Describe  tests  for  (a)  chromium  and  (b)  manganese. 

3.  Review    topics:     (a)  Uses    of    manganese   dioxide.      (b)  Oxidation 
with  potassium  permanganate,     (c)  Special  steels,     (d)  Alums,     (e)  Lead 
chromate. 

4.  Starting  with  MnO2  how  would  you  prepare  in  succession  MnClz 
and  manganese  sulphide?    Starting  with  Fe(CrO2)2  how  would  you  prepare 
K2CrO4,  K2Cr2O7,  potassium  chromate,  lead  chromate? 

5.  Write  the  formulas  of  the  chromate,  dichromate,  manganate  and 
permanganate  corresponding  to  NH4,  calcium,  Pbn,  Al,  magnesium,  Ag, 
Na.     (Use  Valence  Tables.) 

6.  What  is  ferrochrome,  black  oxide  of  manganese,  chrome  yellow, 
chrome  alum,  pyrolusite,  spiegel  iron? 

PROBLEMS 

1.  What  weight  of  the  pure  metals  can  be  prepared  by  the  interaction 
of  aluminium  and  (a)  2  kg.  of  manganese  dioxide,  and  (b)  2000  gm.  of 
chromium  trioxide? 

2.  How  much  potassium  chromate  can  be  made  from  potassium  hy- 
droxide and  200  gm.  of  the  other,  compound? 

3.  What  weight  of  potassium  dichromate  can  be  made  from  3  metric 
tons  of  potassium  chromate? 

4.  Complete  and  balance  the  following:    (a)  CaOO4  +  K2SO4  =  — 

+  K2Cr04;  (6)  Pb(NO3)2  +  -    -  =  PbCrO4  +  KNO3;  (c)  MnO2  +  K2CO3 

+  O  =  K2MnO4  H . 

6.  What  is  the  atomic  weight  of  chromium,  if  6.6595  8m-  °f  ammonium 
dichromate  yield  4.0187  gm.  of  chromium  trioxide?  (Use  exact  atomic 
weights.) 


CHAPTER  XXXII 

PLATINUM 

518.  Occurrence.  —  Platinum  occurs  as  the  chief  in- 
gredient of  an  alloy  called  platinum  ore.     The  associated 
metals  are  ruthenium,   osmium,   iridium,   rhodium,   and 
palladium.     Iron,    gold,    and    copper    are    also    usually 
present.     Only    one    native    compound    is    known,    viz. 
platinum  arsenide  (sperrylite,  PtAs2). 

519.  Preparation.  —  Platinum  is  obtained  as  a  spongy  mass  by 
subjecting  its  alloy  to  a  complicated  treatment  which  involves  boiling 
with  aqua  regia,  then  precipitation  with  ammonium  chloride,  and  final 
heating.     This  spongy  platinum  is  melted  in  a  lime  crucible  with  an 
oxyh'ydrogen  flame,  or  in  an  electric  furnace,  and  hammered  while 
hot  into  a  compact  form. 

520.  Properties   and  Uses.  —  Platinum  is  a  lustrous, 
silvery  metal.     It  is  malleable  and  ductile.     Although  it 
is  attacked  by  fused  caustic  alkalies,  low  melting  metals, 
and  aqua  regia,  it  is  practically  indispensable  in  the  chem- 
ical laboratory,  owing  to  its  high  melting  point  (1755°  C.) 
and   its   resistance   to  many   chemicals.      Platinum  is  a 
good  conductor   of   electricity  and  has  about  the  same 
coefficient  of  expansion  as  glass;  these  properties  adapt 
it  for  use  in  incandescent  electric  light  bulbs.     Electrodes 
are  often  made  of  platinum.     A  rigid  alloy  of  platinum 
is  used    by  dentists  as  a  support  for    teeth.     Recently 
platinum  has  come  into  use  as  jewelry,  especially  in  the 
form  of  mountings  for  gems.     Platinum  has  a  specific 
gravity  of  about  21,  which  is  higher  than  that  of  any 
known  substance,  except  osmium  and  iridium.     In  the 


PLATINUM  381 

form  of  a  black,  porous  mass  it  is  called  spongy  plati- 
num, and  a  still  finer  form  is  called  platinum  black. 
Asbestos  coated  with  platinum  is  used  as  a  catalyzer  in 
manufacturing  sulphuric  acid  by  the  contact  process 
(286,  287).  Platinum  forms  alloys  with  other  metals, 
and  should  never  be  heated  with  lead,  similar  metals, 
or  their  compounds,  since  the  alloys  have  a  low  melting 
point.  With  iridium,  however,  it  forms  a  very  hard 
alloy  of  which  certain  standard  metric  apparatus  is  made. 

521.  Compounds.  —  Chloroplatinic    acid   (H2PtCl6)  is  formed  by 
dissolving  platinum  in  aqua  regia;    it  yields  the  sparingly  soluble, 
yellow    salts  potassium  chloroplatinate    (K2PtCl6)  and    ammonium 
chloroplatinate  (  (NH4)2PtCl6). 

522.  The   metals   associated   with   platinum   have   limited  uses. 
Palladium  is  used  in  chemical  analysis  to  absorb  hydrogen,  and  a 
native  (as  well  as  an  artificial)  alloy  of  iridium  and  osmium,  called 
iridosmine  or  osmiridium,  is  used  to  tip  gold  pens. 

PROBLEMS 

1.  A  piece  of  platinum  foil  measuring  10.5  cm.  by  1.5  cm.  weighs  0.723 
gin.     Into  how  many  pieces,  each  weighing  i  dg.,  may  it  be  divided? 

2.  The  specific  heat  of  platinum  is  0.0324.     According  to  analysis,  35.5 
gm.  of  chlorine  unite  with  48.6  gm.  of  platinum  to  form  platinic  chloride. 
What  is  (a)  the  atomic  weight  of  platinum  and  (b)  the  formula  of  platinic 
chloride? 

3.  Calculate  the  weight  of  platinum  in  (a)  25  gm.  of  potassium  chloro- 
platinate, (b)  i  kg.  of  ammonium  chloroplatinate,  and  (c)  500  mg.  of  barium 
platinocyanide  (BaPt(CN)4). 


CHAPTER  XXXIII 

RADIUM  AND  RADIOACTIVITY 

523.  Occurrence  of  Radium.  —  Radium  is  a  constitu- 
ent of  certain  rare,  uranium-bearing  minerals,  especially 
pitchblende    and    carnotite.     Pitchblende    is    found    in 
Bohemia  and  carnotite  in  Colorado  and  Utah. 

524.  Preparation  of  Radium  Compounds. -- The  pro- 
portion of  radium  in  pitchblende  and  carnotite  is  minute, 
only  a  few  milligrams  to  the  ton.     This  small  proportion 
of  radium,  together  with  considerable  barium,  is  carefully 
extracted  from  these  minerals  by  a  complicated  chemical 
process;    the  radium  is  then  separated  from  the  barium 
as  radium  chloride  or  bromide  by  a  tedious  process  of 
crystallization.     Radium  is  sold  usually  as  radium  bro- 
mide (RaBr2);    the  price  varies  with  the  purity  of  the 
salt,  but  it  is  $100  or  more  a  milligram.     The  supply  is 
exceedingly  limited. 

525.  Properties    of    Metallic    Radium     and    Radium 
Compounds.  —  The  general  properties  of  radium  indicate 
that  it  belongs  to  the  alkaline  earth  family.     Metallic 
radium,  which  was  first  isolated  by  Madame  Curie  in 
1910,  closely  resembles  barium.     Both  are  silvery  white 
metals    and    have    similar    spectra.     Radium    forms    a 
chloride  (RaCl2)  and  a  sulphate  (RaSO4)  whose  properties 
are  like  those  of  the  corresponding  barium  compounds. 
The    bromide    (RaBr2)    is    the    salt    most    often    used; 
indeed  the  actual  work  has  been  largely  done  with  this 
compound  and  not  with  metallic  radium,   despite   the 
fact  that  the  word  radium  is  almost  exclusively  used. 


RADIUM   AND   RADIOACTIVITY  383 

Radium  compounds  color  the  Bunsen  flame  red.  They 
are  self-luminous,  and  a  tube  containing  a  radium  salt 
can  readily  be  seen  in  a  dark  room.  They  cause  fluores- 
cence (i.e.  glowing)  in  various  substances,  e.g.  diamond, 
zinc  sulphide  (ZnS),  willemite  (Zn2SiO4),  and  barium 
platinocyanide  (BaPt(CN)4).  This  fact  is  sometimes 
utilized  to  distinguish  genuine  from  spurious  diamonds. 
Radium  salts  decompose  many  stable  chemical  compounds 
and  cause  chemical  reactions.  Thus,  a  radium  salt 
turns  sodium  glass  brown  and  potassium  glass  purple, 
owing  to  the  liberation  of  sodium  and  potassium  from 
the  glass;  it  transforms  oxygen  into  ozone,  and  yellow 
phosphorus  into  red.  Radium  compounds  sterilize  seeds 
and  kill  microorganisms;  they  also  cause  burns  and 
disintegrate  tissue.  Possibly  radium  preparations  may 
prove  effective  in  curing  certain  skin  diseases  and  malig- 
nant growths. 

526.  Special  Properties  of  Radium  Compounds.  - 
Besides  the  properties  just  mentioned,  radium  compounds 
have  others  which  are  very  characteristic  and  are  not 
exhibited  by  most  substances,  (i)  Radium  compounds 
spontaneously  evolve  considerable  heat;  the  compounds 
are  always  a  little  warmer  than  the  surrounding  air.  It 
has  been  estimated  that  pure  radium  would  liberate 
enough  heat  every  hour  to  raise  its  own  weight  of  water 
from  the  freezing  point  to  the  boiling  point.  (2)  Radium 
compounds  affect  a  photographic  plate  just  as  light 
does.  If  a  tube  containing  a  radium  compound  is  left  a 
short  time  on  a  photographic  plate  (wrapped  in  black 
paper),  or  even  drawn  slowly  across  it,  an  image  is  pro- 
duced when  the  plate  is  developed,  just  as  if  a  photograph 
had  been  taken  in  the  usual  way  (Fig.  98).  (3)  Radium 
compounds  make  the  surrounding  air  a  conductor.  For 


384 


CHEMISTRY 


example,  if  they  are  brought  near  a  charged  body,  such 
as  an  electroscope,  they  discharge  it.  This  is  a  very 
delicate  test  and  is  used  to  detect  radium  compounds  as 

well  as  to  determine  their  pro- 
portion in  mixtures.     The  spe- 
cial properties   exhibited   by 
radium    compounds    are   called 
radioactive  properties;     similar 
ones  are  possessed  by  com- 
pounds  of  other  elements,  e.g. 
Fig.  98.  —  Effect  of  Radium    uranium  and   thorium.     Some- 
Compounds  upon  a  Photo-     times    these    properties    are    in_ 
graphic  Plate.      (This  was 

produced  by  slowly  writing     eluded  by  the  term  radioactivity. 

on  the  plate  with  a   tube          ^    Discovery  of  Radium.  -  About 
containing  radium  bromide  ^    discovered    b      Hend 

and    then    developing    the 
late)  Becquerel    that    uranium    compounds 

affect  a  light-proof  photographic  plate. 

Some  minerals  containing  uranium  compounds,  particularly  pitch- 
blende, were  later  (1898)  found  by  Madame  Curie  to  be  more  radio- 
active than  uranium  itself.  She  studied  pitchblende  carefully  and 
subsequently  in  collaboration  with  her  husband  extracted  from  this 
mineral  a  minute  quantity  of  a  new  substance  which  was  exceedingly 
radioactive.  The  elementary  constituent  in  it  was  named  radium. 
Since  then,  although  very  small  amounts  of  radium  compounds  are 
available,  radioactivity  has  been  zealously  studied  by  Madame  Curie, 
Rutherford,  and  others. 

528.  Interpretation  of  Radioactivity.  —  Many  inter- 
esting experiments  show  that  radioactivity  is  due  to  the 
spontaneous  emission  from  radium  compounds  of  three 
types  of  radiations,  which  are  called  alpha  (a),  beta  (j3), 
and  gamma  (7)  rays.  The  alpha  rays  consist  of  a  stream 
of  positively  charged  particles  moving  with  great  velocity 
-  from  10,000  to  20,000  miles  a  second.  To  the  alpha 
particles  are  ascribed  most  of  the  electrical  effects,  such 


RADIUM    AND    RADIOACTIVITY  385 

as  discharging  an  electroscope.  They  are  regarded  as 
being  identical  with  positively  charged  helium  atoms; 
their  weight  has  been  calculated  to  be  about  four  times 
that  of  a  hydrogen  atom.  The  beta  rays  carry  a  negative 
charge  of  electricity  and  move  with  varying  velocity, 
which  is  sometimes  almost  as  great  as  the  velocity  of 
light  (186,000  miles  a  second).  The  beta  rays  are  the 
most  efficient  in  affecting  a  photographic  plate.  They 
behave  like  the  cathode  rays  developed  in  a  vacuum 
tube,  i.e.  they  are  streams  of  electrons  or  corpuscles  - 
the  subatomic  particles  of  the  physicist.  The  gamma 
rays  have  the  least  electrical  and  photo- 
graphic power,  but  they  are  the  most  pene- 
trating. Gamma  rays  are  not  material 
particles,  but  pulsations  in  the  ether  similar 
to  X  rays. 

A  simple  instrument  called  the  spinthariscope  (Fig. 
99)  shows  in  a  striking  way  that  particles   are  being 
shot  off  continuously  from  a  radium  compound.     The         g' .    /    . 
screen  S  is  coated  with  zinc  sulphide  and  on  the  needle         scope 
R  there  is  a  minute   quantity  of  radium  bromide. 
Upon  looking  into  the  spinthariscope  through  the  lens,  minute  flashes 
of  light  are  seen  on  the  screen.     The  flashes  are  due  to  the  impacts 
of  the  steady  stream  of  alpha  particles  which  fall  upon  the  screen  and 
produce  fluorescence  in  the  zinc  sulphide. 

529.  Disintegration  of  Radium  Compounds. — Although 
radium  is  an  element  which  possesses  many  properties  like 
those  of  the  other  eighty  or  more  elementary  substances, 
it  differs  from  most  of  them  in  being  unstable.  That  is, 
radium  is  slowly  disintegrating.  It  is  very  generally  be- 
lieved that  this  disintegration  is  manifested  not  only  by  the 
unalterable  properties  included  by  the  term  radioactivity, 
but  also  by  the  production  of  radium  from  uranium 
and  the  formation  from  radium  of  a  series  of  -products. 


386  CHEMISTRY 

Uranium,  the  heaviest  of  all  the  elements,  is  regarded  as  the 
parent  substance.  The  others  in  the  series  are  uranium  X, 
ionium,  radium,  niton,  radium  A,  B,  C,  D,  E,  and  F. 
Some  of  these  substances,  all  of  which  are  believed  to  be 
chemical  elements,  are  very  unstable  and  disintegrate 
rapidly.  Helium  is  given  off  in  some  of  the  transitions. 
It  is  not  known  what  the  final  products  of  disintegration 
are;  some  evidence  indicates  that  lead  is  one.  The 
theory  has  been  proposed  that  atoms  of  radium  and 
other  radioactive  elements  are  slowly  disintegrating  into 
simpler  atoms  and  that  the  disintegration  continues  until 
some  more  or  less  stable  form  is  reached,  e.g.  niton  and 
helium.  Many  observations,  among  them  the  experimental 
demonstration  of  the  production  of  helium  from  radium 
by  Ramsay  and  others,  give  this  theory  some  foundation. 

530.  Other  Radioactive  Elements.  —  Besides  uranium 
and  radium,  the  element  thorium  is  radioactive,  though 
to  a  much  less  degree  than  radium.     Much  that  has  been 
said  above  about  radium  compounds  applies  in  general 
to    thorium    compounds.     Actinium    and  polonium   are 
radioactive  elements. 

531.  Conclusion.  —  Radium     and     other    radioactive 
elements  are  being  carefully  investigated.     The  following 
facts  are  well  established:    (i)  uranium,  thorium,  radium, 
and  possibly  other  elements  are  undergoing  spontaneous 
decomposition;    (2)  helium  is  formed  by  the  decomposi- 
tion of   elements   having  a   higher   atomic   weight;     (3) 
enormous  quantities  of  energy  are  liberated  in  radioactive 
transformations.     Doubtless    in    the    immediate    future 
other  facts  will  be  discovered  which  will  enable  us  to 
understand  more  fully  the  structure  of  atoms  and  the 
relation  of  elements  to  each  other. 


RADIUM   AND   RADIOACTIVITY  387 

EXERCISES 

1.  In  what  minerals  does  radium  occur? 

2.  Discuss  the  preparation  of  radium. 

3.  State  the  properties  of  (a)  radium  and  (6)  radium  compounds. 

4.  State  the  special  properties  of  radium  compounds.    In  what  respect 
are  they  striking? 

5.  Essay  topics:    (a)  Discovery  of  radium,     (b)  Madame  Curie,     (c) 
X-rays,     (d)   Uses  of  radium,     (e)   Properties  of  radium.     (/)  Ramsay. 
(g)  Fluorescence. 

6.  Discuss  (a)  alpha  particles  and  (b)  beta  particles. 

7.  Describe  (a)  a  spinthariscope  and  (b)  an  electroscope.     What  does 
each  show  about  radium  compounds? 

8.  Why  was  radium  so  named? 

9.  Discuss  (a)  helium  and  (b)  niton. 

10.  Review  topics:  (a)   Uranium,   thorium,   and  barium  with  special 
reference  to  radium,     (b)  Photography. 

11.  Discuss  the  disintegration  of  radium. 

12.  Review  atoms  and  the  atomic  theory  in  the  light  of  radioactivity. 

PROBLEMS 

1.  Calculate  the  weight  of  radium  in  .001  gm.  of  (a)  radium  bromide, 
(b)  radium  nitrate,  (c)  radium  sulphate. 

2.  Write  the  formulas  of  the  following  compounds  of  radium:    Iodide, 
fluoride,  carbonate,  acid  carbonate,  oxide,  phosphate  (ortho). 

3.  If  2.61099  milligrams  of  radium  bromide  give  2.00988  milligrams  of 
radium  chloride,  what  is  the  atomic  weight  of  radium?     (Use  exact  atomic 
weights  of  Br  and  Cl.) 


APPENDIX 


1.  The  Metric  System  of  weights  and  measures  is  used  in  chem- 
istry. It  is  based  on  the  meter.  This  is  the  unit  of  length,  and  it 
is  a  little  longer  than  a  yard.  Its  exact  length  is  39.37  inches. 
The  unit  of  weight  is  the  gram.  It  is  a  small  weight,  being  only 
about  one  thirtieth  of  an  ounce.  A  five-cent  coin  weighs  approxi- 
mately five  grams.  The  unit  of  volume  is  the  liter.  It  is  slightly 
larger  than  a  quart. 

TABLE  or  THE  METRIC  SYSTEM 


Length 

Weight 

Volume 

Notation 

Kilometer 

Kilogram 

Kiloliter 

1000. 

Hectometer 

Hectogram 

Hectoliter 

IOO. 

Decameter 

Decagram 

Decaliter 

10. 

METER 

GRAM 

LITER 

i. 

Decimeter 

Decigram 

Deciliter 

O.I 

Centimeter 

Centigram 

Centiliter 

O.OI 

Millimeter 

Milligram 

•    Milliliter 

O.OOI 

TABLE  OF  METRIC  EQUIVALENTS 


i  meter 

=  39.37  inches 

i  inch 

=  2.54  centimeters 

i  kilometer 

=  0.62  mile 

i  mile 

=  1.6  kilometers 

i  centimeter 

=  0.39  inch 

i  cubic  inch 

=  16.39  cubic  centi- 

meters 

i  liter 

=  0.908  quart  (dry) 

quart  (liq.) 

=  0.9465  liter 

i  liter 

=  1.056  quart  (liq.) 

pound  (avoir.) 

=  0.4536  kilogram 

i  gram 

=  15.432  grains 

ounce  (avoir.) 

=  28.35  grams 

i  kilogram 

=  2.2  pounds  (avoir.) 

ounce  (troy) 

=  31.1  grams 

i  metric  ton 

=  2204  pounds 

grain  (apoth.) 

=  0.0648  gram 

APPENDIX  389 

TABLE  OF  METRIC  CONVERSION 


To  Change 

Multiply  by 

Inches  to  centimeters  ........ 

2.ZA. 

Centimeters  to  inches  
Cubic  inches  to  cubic  centimeters  
Cubic  centimeters  to  cubic  inches  
Ounces  to  grams  (avoir.)  
Grams  to  ounces  (avoir.)  
Grains  to  grams  
Grams  to  grains  

Q-3937 
16.387 
0.061 

28.35 
0-0353 
0.0648 

I  r  4? 

In  the  case  of  water  the  following  relation  exists:  i  liter,  i  quart, 
and  1000  cubic  centimeters  weigh  approximately  the  same  as  i  kilo- 
gram, 1000  grams,  and  2.2  pounds.  Since  many  liquids  have  about 
the  same  specific  gravity  as  water,  this  general  relation  is  useful,  and 
should  be  learned.  It  is  clear  from  the  relation  just  given  that  i 
cubic  centimeter  of  water  weighs  i  gram  —  a  fact  to  remember,  since 
this  relation  enables  us  to  convert  volume  into  weight,  and  vice  versa. 

The  customary  abbreviations  of  the  common  denominations  are:  — 

meter,  m.  gram,  gm.  milligram,  mg. 

decimeter,  dm.  kilogram,  kg.  or  Kg.  cubic  centimeter,  cc. 

centimeter,  cm.  decigram,  dg.  liter,  1. 

millimeter,  mm.  centigram,  eg.  cubic  decimeter,  cu.  dm. 

The  same  abbreviation  is  used  for  singular  and  plural,  e.g.  i  m., 
4  gm.,  3  cm.,  50  cc. 

2.  The  Thermometer  in  scientific  use  is  the  centigrade.  The 
boiling  point  of  water  on  this  thermometer  is  marked  100,  and  the 
freezing  point  o.  The  equal  spaces  between  these  points  are  called 
degrees.  The  abbreviation  for  centigrade  is  C.,  and  for  degrees  is  °. 
Thus,  the  boiling  point  of  water  is  100°  C.  Degrees  below  zero  are 
always  designated  as  minus,  e.g.  —  12°  C.  means  12  degrees  below 
zero. 


390 


CHEMISTRY 


The  thermometer  in  popular  use  is  the  Fahrenheit.  On  this 
instrument  the  boiling  point  of  water  is  212°  and  the  freezing  point  is 
32  above  zero. 

To  change  Fahrenheit  degrees  into  the  equivalent  centigrade 
degrees,  subtract  32  and  multiply  the  remainder  by  f  ,  or  briefly  — 

C  =*(F-32). 

To  change  centigrade  degrees  into  the  equivalent  Fahrenheit 
degrees,  multiply  by  f  and  add  32  to  the  product,  or  briefly  — 


The  point  —  273°  C.  is  called  absolute  zero.  Absolute  temperature 
is  reckoned  from  this  point.  Degrees  on  the  absolute  scale  are  found 
by  adding  273  to  the  readings  on  the  centigrade  thermometer.  Thus, 
273°  absolute  is  o°  C.,  274°  absolute  is  +  i  C.,  etc. 


PROBLEMS 

1.  Change  into  Fahrenheit  readings  the  following  centigrade  readings: 
(a)  40,  (&)  25,  (c)  87,  (<*)  -20,  (e)  o,  (/)  120,  (g)  862,  (h)  -40. 

2.  Change  into  centigrade  readings  the  following  Fahrenheit  readings: 
(a)  210,  (&)  18,  (c)  o,  (d)  -20,  («)  212,  (/)  70,  (g)  -40,  (h)  127. 

3.  Express   the    following   centigrade   readings   in   absolute   readings: 
(a)  o,  (b)  100,  (c)  -23,  (d)  250. 

3.   Weights  of  Gases.  —  The  weight  in  grams  of  one  liter  of  gases 
at  o°  C.  and  760  mm.  is  — 


Acetylene 

1.162 

Hydrogen  chloride 

1.64 

Air 

1.293 

Hydrogen  sulphide 

1-537 

Ammonia 

•77 

Methane 

.717 

Carbon  dioxide 

1.977 

Nitric  oxide 

1-34 

Carbon  monoxide 

1.25 

Nitrogen 

1.25 

Chlorine 

3.22 

Nitrous  oxide 

1.977 

Ethylene 

1.25 

Oxygen 

1.429 

Hydrogen 

.0898 

Sulphur  dioxide 

2.927 

4.   The  Vapor  Pressure  of  water  vapor  in  millimeters  of  mercury 
is  as  follows:  — 


APPENDIX 


Temperature 

Pressure 

Temperature 

Pressure 

Temperature 

Pressure 

12 

10.46 

17 

14.42 

22 

19.66 

•5 

10.80 

•5 

14.88 

•5 

20.27 

13 

u.  16 

18 

I5-36 

23 

20.89 

•5 

n-53 

•5 

15.85 

-5 

21-53 

14 

11.91 

iQ 

16.35 

24 

22.18 

•5 

12.30 

-5 

16.86 

•5 

22.86 

IS 

12.70 

20 

17-39 

25 

23-55 

•5 

13.11 

-5 

17.94 

-5 

24.26 

16 

13-54 

21 

18.50 

26 

24.99 

•5 

13-97 

•5 

19.07 

•5 

25-74 

The  numbers  in  the  columns  marked  Pressure  are  the  values  for 
a  in  the  formula  for  the  reduction  of  gas  volumes  given  on  page 
41  (this  book). 


INDEX 


Absolute  alcohol,  203. 
Acetic  acid,  204. 

Glacial,  204. 

Test,  207. 
Acetone,  214. 
Acetylene,  172,  391. 

Burner,  ,173. 

Flame,  173. 

From  carbide,  178. 

Generator,  173. 
Acid  calcium  carbonate,  166. 
Acid,  denned,  143. 

Phosphate,  279. 

Reaction,  82. 

Steel,  344. 
Acidity,  145. 
Acids,  79,  81. 

Dibasic,  145. 

Dissociation,  151. 

General  properties,  142. 

Monobasic,  145. 

Names,  84. 

Salts,  145. 

Tests,  82. 

Tribasic,  145. 
Actinium,  386. 
Agate  ware,  246. 
Air,  104,  391. 

A  mixture,  105. 

Composition,  106. 

Liquid,  109. 

Nitric  acid,  94. 

See  Atmosphere. 

Weight  of  liter,  105,  390. 
Albumin,  212. 

And  mercury,  362. 
Albuminoids,  212. 
Alchemists,  80. 
Alcohol,  202. 

Absolute,  203. 

And  water,  43. 

Beverages,  204. 


Denatured,  203. 

Ethyl,  202. 

Fermentation,  203. 

Grain,  202. 

Manufacture,  203. 

Methyl,  202. 

Wood,  202. 
Alizarin,  176. 
Alkali  metals,  288,  300. 
Alkaline  earth  metals,  318. 
Alkaline  reaction,  83. 
Allotropism,  161. 

Carbon,  161. 
Allotropy,  161. 
Alloys,  antimony,  284. 

Chromium,  374. 

Copper,  307. 

Lead,  370. 

Silver,  311. 

Steel,  346. 

Tin,  366. 

Alpha  particles,  384. 
Alum,  333. 

And  water,  37. 

Baking  powder,  334. 

Chromium,  376. 

Iron,  349. 
Alumina,  331. 
Aluminates,  332. 
Aluminium,  328. 

Bronze,  307. 

Chloride,  334. 

Cleaning  by,  311. 

Hydroxide,  332. 

Metallurgy,  328. 

Oxide,  328,  331. 

Properties,  329. 

Sulphate,  333. 

Test,  332. 

Uses,  330. 

Aluminum.     See  Aluminium. 
Alundum,  332. 


INDEX 


393 


Amalgamation  process,  silver,  309. 
Amalgams,  361. 
Ammonia,  89,  391. 

As  refrigerant,  92. 

Composition,  91. 

Formation,  89. 

Liquid,  90,  92. 

Preparation,  89,  90. 

Properties,  90,  91. 

Test,  90,  91. 

Water,  90. 

Ammoniacal  liquor,  89. 
Ammonium,  94,  300. 

As  metal,  94. 

Carbonate,  302. 

Chloride,  91,  300. 

Chloroplatinate,  381. 

Compounds,  93,  94,  300. 

Bichromate,  87. 

Ferric  citrate,  350. 

Hydroxide,  93. 

Nitrate,  301. 

Sulphate,  301. 

Test,  300. 
Amorphous,  carbon,  154,  157. 

Sulphur,  229. 
An  atmosphere,  104. 
Anaesthetic,  172,  214. 

Nitrous  oxide,  101. 
Anhydride,  42,  166. 
Anhydrite,  323. 
Anhydrous  compounds,  46. 
Aniline,  176. 

Animal  charcoal,  159,  160. 
Anions,  137,  149. 

And  anode,  149. 
Anode,  137. 
Anthracene,  176. 
Anthracite  coal,  157. 
Antimony,  283. 

Alloys,  284. 

Compounds,  283,  284. 

Test,  284. 
Apatite,  266. 
Aquafortis,  98. 

Regia,  80. 
Argol,  206. 
Argon,  108,  355. 


And  nitrogen,  105. 

Atom  in  molecule,  123. 

Discovery,  108. 
Aristotle,  104. 
Arrhenius,  137. 
Arsenic,  282. 

Compounds,  282,  283. 

Test,  283. 

White,  282. 
Arsenious  oxide,  282. 
Arsenopyrite,  282. 
Arsine,  283. 
Asbestos,  354. 
Atmosphere,  86,  104. 

Ingredients,  105. 

Rare  gases,  109. 
Atmospheric  pressure,  104,  105. 
Atom  and  radicals,  83. 

And  radioactivity,  58. 

Weight,  60. 
Atomic  theory,  57. 

And  chemical  change,  58. 

And  laws,  58. 
Atomic  weight,  57. 

And  valence,  132. 

Problems,  125. 
Atomic  weights,  60. 

Approximate,  116. 

Determination,  116-120. 

Formula,  62. 

International,  120. 

Symbols,  60. 

Table,  inside  back  cover. 
Atoms,  57. 

And  ions,  137. 

In  molecule,  115,  116,  123. 
Automatic  sprinkler,  285. 
Avogadro's  hypothesis,  114, 115, 117. 
Azurite,  303,  309. 

Babbitt  metal,  370. 
Bacteria,  88. 

In  soil,  88,  98. 
Baking  powder,  206,  29.3. 

Phosphate,  279. 

Tartrate,  206. 
Balancing  equations,  68. 
Balard,  270. 


394 


INDEX 


Barium,  325. 

Compounds,  326. 

Test,  326. 

Barometer,  104,  105. 
Base,  83. 

Denned,  143. 

Diacid,  145. 

Dissociation,  151. 

General  properties,  142. 

Monacid,  145. 

Names,  84. 

Triacid,  145. 
Basic,  salt,  145. 

Steel,  344. 
Basicity,  145. 
Bauxite,  328. 
Beer,  204. 
Beet  sugar,  194. 
Benzaldehyde,  214. 
Benzene,  122,  176. 
Benzine,  122,  175,  176. 
Benzol,  176. 
Berthollet,  56. 
Bessemer  steel,  343. 
Beta  particles,  385. 
Bismuth,  284. 

Alloys,  285. 

Compounds,  286. 

Test,  286. 

Bitter  almonds,  214. 
Bittern,  270. 
Bituminous  coal,  157. 
Bivalent  element,  127. 
Black  damp,  171. 

Lead,  155. 
Blast  furnace,  339. 
Blasting  gelatin,  210. 
Bleaching  by  chlorine,  75,  76. 

Hydrogen  dioxide,  54. 

Sulphur  dioxide,  233. 
Bleaching  powder,  75. 
Blood,  213. 

And  carbon  monoxide,  169. 
Blowpipe,  191. 

Acetylene,  173. 

Flame,  191. 

Oxy-hydrogen,  27. 
Blue  print  paper,  350. 


Blue  stone,  308. 

Bluing,  351. 

Body,  elements  in,  8. 

Boiling  point,  solutions,  139. 

Bomb  calorimeter,  220. 

Bone  black,  160. 

Bones,  275,  280. 

Borax,  245. 

Beads,  246. 

Test,  247. 

Bordeaux  mixture,  308. 
Boric  acid,  246. 
Boron,  245. 

Oxides,  245. 

Test,  247. 
Bort,  155. 
Boyle's  law,  32. 
Brandy,  204. 
Brass,  307. 
Bread,  and  dextrin,  200. 

Composition,  215. 

Making,  199. 
Brimstone,  227. 
Brine,  77,  92. 
Britannia  metal,  367. 
Bromides,  270. 
Bromine,  269. 

Compounds,  270. 

Water,  270. 
Bronze,  307. 
Bunsen,  188. 

'Burner,  188. 

Flame,  188,  189,  190. 
Bureau  of  Mines,  20. 
Burettes,  144. 
Butter,  209. 

Acids,  205. 

Composition,  215. 

Fat,  209. 
Butyric  acid,  205,  209. 

Cadmium,  360. 

Atom  in  molecule,  123. 
Caffeine,  214. 
Calcite,  319. 
Calcium,  318. 

Acid  sulphite,  234. 

Carbide,  178. 


INDEX 


395 


Calcium,  carbonate,  319. 

Chloride,  325. 

Chroma te,  375. 

Compounds,  318,  325. 

Cyanamide,  325. 

Hard  water,  324. 

Hydroxide,  322. 

Light,  28. 

Nitrate,  94,  325. 

Oxalate,  325. 

Oxide,  320. 

Phosphate,  281. 

Preparation,  318. 

Sulphate,  323. 

Sulphide,  325. 

Test,>325. 

Calculations,  chemical,  70. 
Calomel,  362. 
Calorie,  220. 

Coal,  157. 

For  one  dollar,  224. 

Large,  220. 

Small,  144. 
Calorimeter,  220. 
Candle  flame,  186. 

Power,  184,  190. 
Cane  sugar,  194. 
Caramel,  195. 
Carat,  diamond,  155. 

Gold,  315. 
Carbide,  161,  177. 

Calcium,  178. 

Silicon,  177. 
Carbohydrates,  194. 

Digestion,  217. 
Carbolic  acid,  176. 
Carbon,  amorphous,  154,  157. 

Chemical  properties,  161. 

Compounds,  154. 

Cycle,  165. 

Fuel  value,  157,  161. 

Occurrence,  154. 

Test,  1 60. 

Tetrachloride,  77. 
Carbon  dioxide,  390. 

And  life,  164. 

And  plants,  164,  165. 

Formalin,  162. 


In  air,  107. 

Occurrence,  162. 

Preparation,  163. 

Properties,  163. 

Test,  162,  323. 

Vapor  density,  116. 
Carbon  disulphide,  242. 
Carbon  monoxide,  167,  390. 
Carbona,  77. 
Carbonado,  155. 
Carbonates,  166. 
Carbonic  acid,  165. 

Anhydride,  166. 

Oxide,  169. 
Carborundum,  177. 
Carnallite,  296,  354. 
Carnotite,  382. 
Casein,  in  milk,  198. 

Milk,  213. 
Cassiterite,  365. 
Cast  iron,  338. 
Castile  soap,  210. 
Catalysis,  234. 
Cathode,  137. 
Cations,  137,  149. 

And  cathode,  149. 
Caustic  potash,  299. 

Soda,  293. 
Cavendish,  108. 
Celluloid,  201. 
Cellulose,  200. 

Nitrates,  201. 
Cement,  322. 

Portland,  322. 
Cerium,  373. 

Oxide,  190. 
Chalk,  320. 
Charcoal,  158. 

Animal,  159,  160. 

Sugar,  195. 

Wood,  159. 
Charles'  law,  31. 
Cheese,  198,  213. 
Chemical  change,  3. 

And  atomic  theory,  57,  58. 

Combination,  15. 

Decomposition,  13. 

Oxygen,  12,  14. 


396 


INDEX 


Chemical  change,  substitution,  24. 
Chemical  properties,  4. 

Reaction,  65. 
Chemistry,  i. 
Chile  saltpeter,  295. 
Chinese  white,  359. 
Chloride  of  lime,  76. 
Chlorides,  75,  77,  So. 

Names,  81. 

Test,  81. 

Chlorination  process,  314. 
Chlorine,  73,  390. 

And  ammonia,  91. 

Atomic  weight,  119. 

Manufacture,  294. 

Molecular  formula,  123. 

Preparation,  73. 

Properties,  74,  75. 

Water,  74. 
Chloroform,  172. 
Chlorophyl,  165,  337. 
Chocolate,  214. 
Choke  damp,  171. 
Chrome  iron  ore,  374. 
Chromic  acid,  377. 
Chromite,  374. 
Chromium,  374. 

Alum,  376. 

Compounds,  373,  374,  376. 

Preparation,  330. 

Steel,  374. 

Test,  375,  376. 
Cider,  205. 
Cinnabar,  360,  363. 
Citric  acid,  207. 
Clay,  328,  335. 
Coal,  157. 

Anthracite,  157. 

Bituminous,  157. 

Calorific  value,  157. 

Composition,  157. 

Fire,  168. 

Formation,  158. 

Lignite,  157. 

Soft,  89. 

Tar,  176,  182. 
Coal  gas,  181. 

Manufacture,  181. 


Cobalt,  351. 

Compounds,  351,  352. 

Test,  352. 
Cocoa,  214. 
Coffee,  214. 
Coke,  1 60,  176. 

Ovens,  90. 
Cold  storage,  92,  93. 
Colemanite,  245. 
Collagen,  212. 
Collodion,  201. 
Colloids,  silicic  acid,  254. 
Combination,  15. 

Valence,  129. 
Combining  number,  133. 

Weight,  133. 
Combustion,  16,  162. 

Spontaneous,  16. 
Common  salt,  290. 

Deliquescence,  47. 
Composition,  air,  106. 

Ammonia,  91. 

Constant,  59. 

Gases,  113. 

Hydrogen  chloride,  80. 

Nitric  acid,  97. 

Nitric  oxide,  102. 

Nitrous  oxide,  101. 

Percentage,  62,  63. 

Water,  23. 

Compound,  defined,  5. 
Concentration,  ore,  304. 
Concrete,  322. 
Condenser,  water,  37,  38. 
Conservation,  energy,  221. 

Matter,  55,  58. 
Constant  composition,  55,  59. 
Converter,  copper,  305. 

Iron,  343. 
Cooking  food,  224. 
Copper,  303. 

Alloys,  307. 

And  nitric  acid,  98. 

And  sulphuric  acid,  232,  240. 

Blister,  305. 

Compounds,  303,  307,  308. 

Electrolytic,  305. 

Metallurgy,  303. 


INDEX 


397 


Copper,  nitrate,  98,  308. 

Ores,  303. 

Oxides,  308. 

Properties,  306. 

Refining,  305. 

Sulphate,  308. 

Sulphide,  308. 

Tests,  306. 

Uses,  307. 

Copper  nitrate,  98,  308. 
Copper  sulphate,  308. 

Anhydrous,  308. 

Electrolysis,  149,  305. 

Fehling's  solution,  197. 

Hydrolysis,  147. 
Copperas,  348. 
Coquina,  320. 
Cordite,  210. 
Corpuscles,  385. 
Corrosive  sublimate,  362. 
Corundum,  331. 
Cosmetics,  286. 
Cotton  seed  oil,  209,  210. 
Courtois,  271. 
Cream  of  tartar,  206. 
Crocus,  347. 
Cryolite,  266,  328. 
Crystallization,  44. 
Crystals,  44. 
Cullinan  diamond,  155. 
Cupric  sulphate,  308. 
Cuprous  oxide,  197. 
Curie,  Madame,  384. 
Cyanide  process,  314. 
Cyanogen,  179. 

Compounds,  350. 

Dal  ton,  57. 

Davy,  74,  101,  171,  272,  288,  296. 

Decomposition,  13. 

Double,  65. 
Decrepitation,  290. 
Deflagration,  100. 
Deliquescence,  47. 
Denatured  alcohol,  203. 
Desiccator,  47. 
Dew  point,  106. 
Dewar  flask,  no,  197. 


Dextrin,  200. 
Dextrose,  195. 

Reducing  by,  196. 

Test,  197. 

Diamond,  154,  177,  383. 
Diastase,  198. 
Diatoms,  249. 
Dibasic  acid,  145. 
Diffusion,  24. 
Disinfectant,  214. 
Displacement,  24. 
Dissociation,  electrolytic,  137. 

Acids,  bases,  salts,  151. 
Distillation,  destructive,  160. 

Fractional,  175. 

Water,  37. 

Divalent  element,  127. 
Dolomite,  354,  357. 
Double  decomposition,  65. 
Dulong  and  Petit,  119. 
Dyad,  127,  129. 
Dyeing,  334. 
Dynamite,  210. 

Earthenware,  336. 
Earth's  crust,  8. 
Effervescence,  42. 
Efflorescence,  46,  47. 

EggS,   212. 

Decayed,  229. 

Preserving,  253. 
Electric  furnace,  177,  178. 

Calcium  carbide,  178. 

Carbon  disulphide,  243. 

Phosphorus,  276. 
Electricity  and  solutions,  136. 
Electrodes,  137. 

And  ions,  149. 
Electrolysis,  148,  150. 

Aluminium  oxide,  328. 

Calcium  chloride,  318. 

Carnallite,  354. 

Copper  sulphate,  149,  305. 

Gold,  315. 

Illustrations,  148. 

Interpretation,  148. 

Lead,  369. 

Sodium  chloride,  73,  293. 


398 


INDEX 


Electrolytes,  136. 

And  electrolysis,  148. 

Boiling  point,  139. 

Chemical  behavior,  140. 

Freezing  point,  139. 
Electrolytic  cell,  148. 
Electrolytic  dissociation,  137. 

Acids,  bases,  salts,  137,  151. 

And  salts,  146. 

Facts,  139. 
Electrons,  385. 
Electroplating,  151. 
Electro-silicon,  251. 
Electrotyping,  150,  156. 
Element,  6. 

And  atoms,  57. 
Elements,  classification,  259. 

Distribution,  7,  8. 

In  body,  8,  218. 

Inert,  87. 

Molecular  formulas,  123. 

Molecular  weights,  123. 

Radioactive,  386. 

Table  7,  inside  back  cover. 
Emery,  331. 
Energy,  conservation,  221. 

From  food,  218. 
Enzymes,  196,  216. 
Epsom  salts,  356. 
Equations,  65-70,  123,  124. 

Ionic,  141. 

Thermal,  161. 

Volumetric,  124. 
Equivalents,  132-134. 
Esters,  207. 

Fats,  208 
Etching,  268. 
Ether,  214. 

And  water,  43. 
Ethyl,  193. 

Acetate,  207. 

Alcohol,  202,  203,  207. 

Ether,  214. 
Ethylene,  172,  391. 
Explosions,  coal  mine,  171. 

Factors,  66. 
Fat, -208. 


Fat,  digestion,  217. 
Fehling's  solution,  197. 
Fermentation,  163,  196. 

Acetic,  205. 

Alcoholic,  196,  203. 

Lactic,  198. 

Maltose,  198. 

Sugar,  196. 
Ferric  chloride,  349. 

Compounds,  347. 

Ferrocyanide,  350. 

Hydroxide,  348. 

Oxide,  347. 

Sulphate,  348. 

Sulphide,  349. 

Sulphocyanate,  351. 
Ferrite,  iron,  348. 
Ferrochrome,  346,  374. 
Ferrosilicon,  346. 
Ferrous  compounds,  347. 

Carbonate,  349. 

Chloride,  349. 

Ferricyanide,  350. 

Hydroxide,  348. 

Oxide,  347. 

Sulphate,  348. 

Sulphide,  349. 
Fertilizer,  nitrogen,  88,  94. 

Calcium  nitrate,  94. 

Cyanamide,  325. 

'Phosphate,  281. 

Potassium,  300. 

Sodium  nitrate,  295. 
Fertilizers,  calcium,  325. 
Filter,  water,  37. 
Fire  damp,  1 70. 
Fire  extinguisher,  77,  164. 
Fireproof  door,  285. 
Fireworks,  298,  325,  326. 
Fixation  of  nitrogen,  88. 
Flame,  185.  » 

Acetylene,  173. 

Bunsen,  188. 

Candle,  186. 

Cooled,  1 88. 

Hydrogen,  25,  27,  188. 

Luminous,  186,  187,  188 

Nature,  185. 


INDEX 


399 


Flame,  non-luminous,  188. 

Ordinary  gas,  187. 

Oxidizing,  190,  191. 

Oxy-acetylene,  18,  173. 

Oxy-hydrogen,  18. 

Reducing,  190,  191. 

Reversed,  185. 

Structure,  186. 
Flavoring  extracts,  207. 
Flavors,  214. 
Flour,  199. 

Proteins,  212. 
Fluorescence,  383. 
Fluorides,  255,  268. 
Fluorine,  266. 
Fluorite,  266. 
Fluor  spar,  266. 
Fluosilicic  acid,  255. 
Food,  214. 

And  energy,  218. 

Composition,  215. 

Dietary  studies,  223. 

Fuel  value,  219,  221. 

Nitrogenous,  88. 

Nutritive  ratio,  222. 

Nutritive  value,  222. 

Relative  cost,  224. 

Table,  216,  221,  223,  224. 
Foodstuffs,  215. 
Fool's  gold,  349. 
Formaldehyde,  214. 
Formalin,  214. 
Formula,  60. 

Calculation,  62. 

Composition,  62,  63. 

From  valence,  130. 

Graphic,  132. 

Molecular,  122. 

Molecular  weight,  61. 

Simplest,  62,  121,  122. 

Structural,  132. 

Writing,  130. 
Free  alkali,  211. 
Freezing  point,  139. 

Solutions,  139. 
Fructose,  196. 
Fruit  sugar,  196. 
Fuel  value  of  food,  219. 


Fuel  value  of  food,  table,  221,  223. 
Fusible  link,  285. 
Fusible  metals,  285, 

Galena,  368,  372. 
Galvanized  iron,  358. 
Gamma  particles,  385. 
Gas,  air,  104. 

Carbon,  160. 

Liquor,  89. 

Mantles,  185. 

Range,  26. 

Gases,  weight  of  liter,  390. 
Gasoline,  175. 

Engine,  175. 
Gay-Lussac,  114. 

Law,  114,  115. 

Tower,  236,  237. 
Gelatin,  212. 

Blasting,  210. 
Gems,  332. 
German  silver,  307. 
Geyserite,  254. 
Glacial  phosphoric  acid,  278. 
Glass,  255. 

Annealing,  257. 

Etching,  268. 

Fluorescent,  377. 

Ingredients,  256. 

Kinds,  256. 

Water,  253. 
Glauber's  salt,  294. 
Globulins,  212. 
Glover  tower,  236,  237. 
Glucose,  196. 
Gluten,  212. 
Glutelins,  212. 
Glycerin,  208,  209. 
Glycerol,  210. 
Glyceryl,  208. 
Glycogen,  200. 
Gold,  313. 

Coin,  307. 

Compounds,  315. 

Cyanide,  316. 

Fool's,  349. 

Metallurgy,  314. 

Noble  metal,  80. 


400 


INDEX 


Gold,  plating,  316. 

Test,  316. 
Grape  juice,  206. 
Grape  sugar,  196. 
Graphite,  155. 

Manufactured,  156. 
Green  fire,  326. 
Groups,  periodic,  262. 
Guano,  281. 
Gun  cotton,  201. 
Gun  metal,  307. 
Gunpowder,  201,  297. 
Gypsum,  323. 

Halogen  family,  266,  273. 
Hard  water,  167,  324. 

Magnesium,  356. 
Hardness  of  water,  324. 
Heat  of  neutralization,  144. 
Helium,  109. 

And  radium,  386. 
Helmet,  oxygen,  19. 
Hemaglobin,  213. 
Hematin,  213. 
Hematite,  337. 
Hexad,  127. 

Hexavalent  element,  127. 
Hydrates,  46. 
Hydriodic  acid,  273. 
Hydrobromic  acid,  270. 
Hydrocarbons,  170. 
Hydrochloric  acid,  77. 

Commercial,  78. 

Electrolysis,  148. 

Preparation,  77. 

Properties,  78,  79. 

Test,  81,  90,  91,  142. 
Hydrocyanic  acid,  179. 
Hydrofluoric  acid,  267. 
Hydrofluosilicic  acid,  255. 
Hydrogen,  22,  390. 

And  electrolysis,  149,  150. 

And  nitric  acid,  98. 

Burning,  25,  26. 

Explosion,  26. 

Flame,  25,  27,  75. 

From  sodium  hydroxide,  289. 

Ions,  143. 


Molecular  formula,  123. 

Name,  24. 

Occurrence,  28. 

Preparation,  22,  23,  24. 

Properties,  24,  25. 

Reduction  by,  26. 

Replacing,  145. 

Test,  27. 

Uses,  27. 
Hydrogen  chloride,  77,  78,  390. 

Composition,  80. 

Formula,  120. 

Test,  79,  90,  91. 
Hydrogen  dioxide,  54. 

Formula,  121. 
Hydrogen  fluoride,  267. 
Hydrogen  sulphide,  229,  390. 

Test,  231. 

Water,  230. 
Hydrolysis,  147. 

Aluminium,  333. 

Antimony,  284. 

Bismuth,  286. 

Copper  sulphate,  308. 

Maltose,  198. 

Sodium  carbonate,  292. 

Sugar,  195. 
Hydroquinone,  313. 
Hydrosulphuric  acid,  230. 
Hydroxides,  83. 
Hydroxyl,  83. 

Groups,  145. 

Ions,  143. 
Hyposulphite,  313. 

Ice,  38. 

Making,  92,  93. 
Iceland  spar,  319. 
Illuminating  gas,  181. 

Candle  power,  184. 

Composition,  184. 
Infusorial  earth,  249,  251. 
Ingots,  344. 
Ink,  indelible,  312. 

Spots,  205. 

Writing,  348. 
Insecticides,  283. 

Sulphur,  229. 


INDEX 


401 


Invertase,  196. 
Iodides,  273. 
Iodine,  271. 

Compounds,  273. 

Molecular  weight,  123. 

Test,  272. 
lodoform,  172. 
Ionic  equation,  141. 
Ionium,  386. 
lonization,  151. 

Per  cent,  151. 

Table,  151. 
Ions,  137,  138. 

And  atoms,  137. 

Kinds,  137. 

Migration,  149. 

Representation,  138. 

Simple,  152. 

Solutions,  140. 

Table,  151,  152. 
Iridium,  381. 
Iron,  337. 

Alum,  349. 

Cast,  338,  340. 

Chlorides,  349. 

Compounds,  337,  347- 

Cyanides,  349. 

Ferrite,  348. 

Galvanized,  358. 

Hydroxides,  348. 

Metallurgy,  338. 

Ores,  335. 

Oxides,  347. 

Pig,  340. 

Pure,  337. 

Russia,  348. 

Rust,  338. 

Stains,  205. 

Sulphates,  348. 

Sulphides,  349. 

Tests,  350. 

Varieties,  338. 

Wrought,  341. 
Isomerism,  197. 
Ivory  black,  160. 

Jelly  making,  202. 
Junket,  213. 


Kainite,  296. 
Kaolin,  335. 
Keratins,  212. 
Kerosene,  175. 

Flashing  point,  175. 
Kieserite,  354,  356. 
Kilogram,  388,  389. 
Kindling  temperature,  17. 
Kipp  apparatus,  22. 
Krypton,  109. 

Lactic  acid,  198,  206,  213. 
Lactose,  197. 
Lampblack,  161. 
Lard,  208,  210. 
Laughing  gas,  101. 
Lavoisier,  17,  23,  87. 
Law,  55. 

Boyle,  32. 

Charles,  31. 

Conservation  of  energy,  221. 

Conservation  of  matter,  55,  58. 

Constant  composition,  55,  59. 

Definite  proportions,  56. 

Dulong  and  Petit,  119. 

Gay-Lussac,  114,  115. 

Multiple  proportions,  56,  59. 

Periodic,  263,  264. 

Specific  heat,  119. 
Lead,  367. 

Acetate,  372,  373. 

Alloys,  370. 

And  radium,  386. 

Arsenate,  283. 

Black,  155,  368,  369. 

Carbonate,  371. 

Chloride,  372. 

Chromate,  373,  374,  376. 

Compounds,  368,  369,  372. 

Dioxide,  371. 

Fluosilicate,  369. 

Metallurgy,  368. 

Monoxide,  370. 

Nitrate,  372. 

Oxides,  370. 

Pencils,  156. 

Properties,  369. 

Sulphate,  372. 


402 


INDEX 


Lead,  tests,  372,  373. 

Tetroxide,  370. 

Uses,  370. 

White,  371. 
Le  Blanc  process,  291. 
Legumes  and  nitrogen,  88. 
Lemonade,  207. 
Levulose,  196. 
Life  and  nitrogen,  88. 

And  carbon  dioxide,  164. 

And  oxygen,  18. 
Lignite,  157. 
Lime,  320. 

And  water,  41. 

Hydraulic,  322. 

Light,  28. 

Milk  of,  323. 

Phosphate,  281. 
Limekiln,  321. 
Limestone,  319,  320. 

Solution,  166,  167. 
Limewater,  323. 
Limonite,  337. 
Liquid  air,  109. 
Liter  of  gases,  weight,  390. 
Litharge,  370. 
Litmus,  82,  83. 

And  salts,  146. 
Loadstone,  348. 
Lockyer,  109. 
Lubricating  oil,  175. 

Magnalium,  331. 
Magnesia,  356,  357. 

Alba,  357. 
Magnesite,  357. 
Magnesium,  354. 

Bromide,  269. 

Carbonate,  354,  357- 

Chloride,  356. 

Compounds,  354. 

Hard  water,  324. 

Hydroxide,  356. 

Nitride,  87,  91,  355. 

Oxide,  355,  356. 

Sulphate,  356. 

Test,  357. 
Magnetite,  337,  348. 


Malachite,  303,  309. 
Malic  acid,  206. 
Malt,  198. 
Maltose,  198. 
Manganese,  377. 

Compounds,  377,  378. 

Dioxide,  377. 

Steel,  346. 

Test,  378. 
Marble,  319. 
Marsh  gas,  170. 
Matches,  14,  280,  298. 
Matte,  304. 
Mendelejeff,  264. 
Mercuric  compounds,  362,  363. 
Mercurous  compounds,  362,  363. 
Mercury,  360. 

Alloys,  361. 

Atom  in  molecule,  123. 

Chlorides,  362. 

Compounds,  360,  362. 

Nitrates,  363. 

Oxide,  n,  12,  362. 

Poisoning,  362. 

Properties,  361. 

Tests,  362,  367. 

Uses,  361. 
Metabolism,  216. 
Metal  and  non-metal,  259,  260. 
Metalloids,  259. 
Metals,  259,  260. 

Alkali,  288,  300. 

Alkaline  earth,  318. 

Displacement,  316. 

Electrothermal,  316. 

Fusible,  285. 

Metaphosphoric  acid,  278. 
Metasilicic  acid,  252. 
Methane,  170,  391. 
Methyl,  193. 
Methyl  alcohol,  202. 
Metric  system,  388. 
Mexican  onyx,  320. 
Migration  of  ions,  149. 
Milk,  197,  209. 

Casein,  213. 

Composition,  215. 

Sour,  206. 


INDEX 


403 


Mineral  matter  in  body,  218. 
Mineral  oil,  176. 
Minium,  370. 
Mispickel,  282. 
Mixture,  5. 

Air,  105. 

Moissan,  155,  266. 
Molasses,  195,  205. 
Molecular  equations,  123. 
Molecular  formula,  122. 

Elements,  123. 
Molecular  weights,  61. 

And  formulas,  122. 

Approximate,  115,  117. 

By  boiling  point,  140. 

By  freezing  points,  140. 

Determination,  115. 

Elements,  123. 
Molecule,  57. 

Formulas,  120. 

Of  compounds,  120. 

Relative  weights,  114 

Solutions,  140. 
Molybdenum,  346,  377 
Monad,  127,  129. 
Monazite  sand,  373. 
Monobasic  acid,  145. 
Mordants,  334. 
Mortar,  323. 
Moth  balls,  176. 
Mucilage,  200. 
Multiple  proportions,  56,  59. 
Muriatic  acid,  77. 

Naphtha,  175. 
Naphthalene,  176 
Nascent  state,  76. 
Natural  gas,  1 70. 
Negative  electrode.  137. 
Neon,  109,  123. 
Neutralization,  83. 

Equation,  143,  144. 

Heat  of,  144. 

Ionic  theory,  143. 

Water,  143. 
Neutral  reaction,  82. 
Newton's  metal,  285. 
Nickel,  351. 


Alloys,  374. 

Coin,  307. 

Compounds,  351. 

Steel,  346. 

Test,  351. 
Nichrome,  374. 
Niter,  296. 
Niton,  386. 
Nitrates,  98. 

Properties,  99,  100,  102. 

Test,  100. 
Nitric  acid,  94,  95,  96. 

Commercial,  95,  96. 

Composition,  97. 

Formation,  94. 

From  air,  94. 

Fuming,  102. 

Manufacture,  95. 

Oxidizing  agent,  97. 

Preparation,  94,  95. 

Test,  100. 
Nitric  oxide,  101,  390. 

Composition,  102. 

Formation,  98,  99. 
Nitrides,  87. 
Nitrification,  98. 
Nitrites,  100. 

Test,  loo. 
Nitrobenzene,  176. 
Nitrogen,  86,  390. 

And  life,  88. 

Animal  matter,  89. 

Discovery,  87. 

Fixation,  88. 

From  liquid  air,  in. 

In  air,  105,  106. 

Molecular  formula,  123. 

Occurrence,  86. 

Oxides,  100. 

Preparation,  86,  91. 

Properties,  87. 
Nitrogen  dioxide,  102. 

Formation,  99. 
Nitrogen  tetroxide,  102. 
Nitroglycerin,  210,  251. 
Nitrose  acid,  237. 
Nitrosyl-sulphuric  acid,  235. 
Nitrous  acid,  100. 


404 


INDEX 


Nitrous  oxide,  100,  302,  390. 

Composition,  101. 
Non-electrolytes,  136. 

Boiling  point,  139. 

Freezing  point,  139. 
Non-metals,  259,  260. 
Normal  pressure,  104. 
Normal  salts,  145. 
Nutrients,  215. 

Comparative  cost,  224. 

Table,  223. 
Nutrition,  214,  215. 

Occlusion,  24. 

Ocean,  composition,  8. 

Water,  36. 
Oils,  mineral,  176. 

Natural,  208. 
Oleic  acid,  205. 
Olein,  208. 
Oleomargarin,  209. 
Olive  oil,  205,  208,  210. 
Opal,  249. 

Open-hearth  process,  344. 
Orange  mineral,  371. 
Ores,  303. 

Treatment,  304. 
Organic  compounds,  193. 
Orpiment,  282. 
Orthophosphoric  acid,  278. 
Orthosilicic  acids,  252. 
Osmiridium,  381. 
Osmium,  381. 
Oxalic  acid,  206. 
Oxidation,  15,  76,  367. 

And  combustion,  16. 

And  reduction,  26,  27. 

Blood,  213. 

Flame,  190. 

In  body,  162. 
Oxides,  15. 
Oxidizing  agent,  15. 

Aqua  regia,  97. 

Chromic  acid,  377. 

Fuming  nitric  acid,  102. 

Nitric  acid,  96. 
Oxidizing  flame,  190,  191. 
Oxygen,  n,  390. 


And  blood,  213. 

And  electrolysis,  150. 

And  life,  18,  213. 

And  vapor  density,  115. 

Atoms  in  molecule,  115,  116. 

Cycle,  165. 

Equivalent  weight,  132. 

From  sodium  hydroxide,  289. 

Helmet,  19. 

In  air,  105,  106. 

In  liquid  air,  no,  in. 

In  water,  74. 

Molecular  formula,  123. 

Molecule,  115. 

Name,  17. 

Occurrence,  n. 

Preparation,  n. 

Properties,  13,  14. 

Test,  14. 

Uses,  1 8. 

Weight  of  liter,  33,  390. 
Oxy-acetylene  flame,  173. 
Oxy-hydrogen,  blowpipe,  27. 

Flame,  27. 
Ozone,  20. 

Paint,  326,  371. 

Palladium,  24,  381. 

Palmitin,  208. 

Palm  oil,  210. 

Paper,  201. 

Paraffin,  176. 

Paris  green,  283. 

Parkes  process,  310. 

Pearlash,  299. 

Pectic  acid,  202. 

Pectin,  202. 

Pectocellulose,  202. 

Pentad,  127,  129. 

Pentavalent  element,  127. 

Percentage  composition,  62,  63. 

Periodic  classification,  260. 

Families,  262. 

Halogens,  273. 

Law,  263. 

Table,  261. 

Permanent  hardness,  324. 
Petroleum,  174. 


INDEX 


405 


Petroleum,  ether,  175. 
Pewter,  367. 
Phenol,  176. 
Phlogiston,  17. 
Phosphate  281. 

Rock,  281. 

Test,  279. 

Phosphates,  275,  279,  280,  281. 
Phosphoproteins,  213. 
Phosphoric  acid,  278,  281. 
Phosphorus,  275. 

And  air,  86. 

And  life,  218,  280. 

And  oxygen,  14. 

Compounds,  279. 

Fertilizer,  281. 

Matches,  280. 

Molecular  weight,  123. 

Ordinary,  276. 

Oxides,  277. 

Pentoxide,  86. 

Red,  277. 

Yellow,  276. 
Photography,  242,  312,  350. 

Radium,  383. 
Physical  change,  2. 
Physical  properties,  4. 
Picromerite,  296. 
Pig  iron,  340. 
Pins,  366. 
Pintsch  gas,  181. 
Pitchblende,  382. 
Plants,  and  nitrogen,  88. 

And  carbon  dioxide,  163. 
Plaster,  323. 

Of  Paris,  323. 
Platinum,  380. 

Compounds,  381. 

Metals,  380,  381, 
Plumbago,  155. 
Polariscope,  197. 
Polonium,  386. 
Porcelain,  335. 
Positive  electrode,  137. 
Portland  cement,  322. 
Potash,  299. 

Caustic,  299. 
Potassium,  295. 


And  life,  299. 

And  water,  23. 

Atom  in  molecule,  123. 

Carbonate,  298. 

Chlorate,  298. 

Chloride,  296. 

Chloroplatinate,  381. 

Chromate,  374,  375. 

Cyanide,  179,  299. 

Dichromate,  374,  375. 

Ferricyanide,  350. 

Ferrocyanide,  350. 

Hydroxide,  299. 

Manganate,  378. 

Nitrate,  296. 

Nitrite,  100,  297. 

Permanganate,  378. 

Sulphate,  299. 

Test,  296. 
Precipitate,  81. 
Precipitation,  81. 
Pressure,  normal,  30. 

Standard,  30. 
Priestley,  n. 
Producer  gas,  168. 
Products,  equation,  67. 
Properties,  i,  4. 
Protamins,  212. 
Proteid,  211. 
Protein,  211. 

Composition,  211. 

Digestion,  217. 

Function,  217. 

Groups,  211. 

Nitrogen  in,  86. 

Plants,  98. 
Proust,  55. 
Prussian  blue,  350. 
Prussic  acid,  179. 
Puddling,  341. 
Pulmotor,  19. 

And  carbon  monoxide,  169. 
Purple  of  Cassius,  316. 
Pyrene  fire  extinguisher,  77. 
Pyroligneous  acid,  204. 
Pyrophosphoric  acid,  278. 
Pyrosulphuric  acid,  242. 


406 


INDEX 


Quadrivalent  element,  127. 
Quartation,  314. 
Quartz,  249. 

Vessels,  250,  251. 
Quicklime,  322. 
Quicksilver,  360. 
Quinquivalent  element,  127. 

Radical,  83. 

Ammonium,  94. 

Organic,  193. 

Valence,  128. 
Radioactivity,  383,  384,  385. 

And  atoms,  58. 
Radium,  382. 

Compounds,  382. 

Discovery,  384. 

Disintegration,  385. 

Products,  386. 
Ramsay,  108. 
Rayleigh,  108. 
Reaction,  65,  69. 

Acid,  82. 

Alkaline,  83. 

Basic,  83. 

Neutral,  82. 
Realgar,  282. 
Red  fire,  325. 
Red  lead,  370. 
Reducing  agent,  26. 
Reducing  flame,.  190,  191. 
Reduction,  26,  27,  367. 

Carbon,  162. 

Carbon  monoxide,  169. 

Flame,  190. 

Relative  humidity,  106. 
Rennet,  198,  213. 
Rennin,  213. 
Replacement,  24,  130. 
Respiration,  18. 

Calorimeter,  220. 
Reverberatory  furnace,  341. 
Richards,  56,  119. 
Rocks,  252. 
Rose's  metal,  285. 
Rouge,  347. 
Rum,  204. 
Rusting,  iron,  338. 


Rusting  tinware,  366. 
Rutherford,  87. 

Saccharose,  194. 
Safety  lamp,  171. 
Sal  ammoniac,  301. 
Saleratus,  293. 
Sal  soda,  291. 
Salt,  common,  290. 

Dairy,  290. 

Glauber's,  294. 
Saltpeter,  297. 

Chile,  271,  295. 
Salts,  82. 

Acid,  145. 

Basic,  145. 

Dissociation,  151. 

Formation,  145. 

General  properties,  142. 

Names,  84. 

Normal,  145. 

Preparation,  97,  98. 

Properties,  145. 

Varied  properties,  146. 
Sand,  248,  251. 

Glass,  255. 

See  Silica. 

Seidlitz  powder,  206. 
Selenite,  323. 
Shot,  282. 
Siderite,  337. 
Silica,  249. 

And  life,  251. 

Yellowstone  Park,  254. 
Silicates,  251. 
Silicic  acid,  251. 

Colloidal,  254. 
Silicides,  248. 
Silicon,  248. 

Carbide,  177. 

Test,  255.  ^ 

Tetrafluoride,  250,  255,  268. 
Silicon  dioxide,  249. 

Properties,  250. 

Uses,  251. 
Silver,  309. 

Alloys,  311. 

And  dextrose,  196. 


INDEX 


407 


Silver,  bromide,  312. 

Chloride,  312. 

Cleaning,  311. 

Coin,  307. 

Compounds,  309,  312. 

Iodide,  312. 

Metallurgy,  309. 

Nitrate,  312. 

Oxidized,  311. 

Plating,  311. 

Properties,  310. 

Sterling,  311. 

Sulphide,  231. 

Test,  312. 

Silverware,  blackening,  231. 
Simplest  formula,  121,  122. 
Siphon,  163. 
Sirup,  195. 
Slag,  304,  338. 
Slaked  lime,  321. 
Smelting,  304. 

See  Metallurgy. 
Smokeless  powder,  201. 
Snow  crystals,  39. 
Soap,  208,  210. 

And  hard  water,  210,  211 

Castile,  210. 

Cleaning  action,  211. 

Floating,  211. 

Manufacture,  211. 
Soda,  292. 

Ash,  291. 

Baking,  292. 

Caustic,  293. 

Cooking,  292. 

Crystals,  291. 

Washing,  292. 
Sodium,  288. 

Acid  sulphite,  234. 

And  water,  22,  289. 

Atom  in  molecule,  123. 

Bicarbonate,  206,  292. 

Carbonate,  147,  291. 

Chloride,  290. 

Compounds,  295. 

Cyanide,  179. 

Dioxide,  295. 

Hydroxide,  288,  293. 


Hyposulphite,  242. 

lodate,  271. 

Manganate,  378. 

Manufacture,  288. 

Nitrate,  295. 

Nitrite,  100. 

Peroxide,  295. 

Silicate,  252,  253. 

Sulphate,  294. 

Test,  289. 

Thiosulphate,  242,  313. 
Sodium  chloride,  290. 

Electrolysis,  73,  293. 
Soft  water,  324. 
Solder,  367. 

Borax,  246. 
Solubility,  denned,  44. 

And  temperature,  43. 

Curve,  44,  45. 

Mutual,  43. 

.Table,  44. 

Terms,  44. 
Soluble  glass,  253. 
Solute,  42. 
Solution,  42. 

And  crystallization,  44. 

Boiling  point,  139. 

Chemical  behavior,  140. 

Colloidal,  254. 

Denned,  5. 

Electric  current,  136. 

Freezing  point,  139. 

Gases,  42. 

Liquids,  43. 

Saturated,  43. 

Solids,  43. 

Supersaturated,  45. 

Water  of  crystallization,  46. 
Solvay  process,  291. 
Solvent,  42. 

Sour  milk,  cooking,  206. 
Specific  heat,  119. 
Speculum  metal,  307,  367. 
Spelter,  358. 
Sperrylite,  380. 
•Spiegel  iron,  344,  377. 
Spinthariscope,  385. 
Spontaneous  combustion,  16. 


408 


INDEX 


Stalactite,  320. 
Stalagmite,  320. 
Standard  conditions,  30. 

Formula,  33. 
Standard  dietary,  222. 
Standard  pressure,  104. 
Stannic  compounds,  367. 
Stannous  compounds,  367. 
Starch,  198. 

And  alcohol,  203. 

Animal,  200. 

Baking  powder,  206. 

Test,  199,  272. 
Starching  clothes,  200. 
Stas,  56. 
Stassfurt  salts,  354. 

Deposits,  296. 
Steam,  39. 

Stearic  acid,  fats,  208. 
Stearin,  208. 
Steel,  342. 

Alloys,  346. 

And  aluminium,  330. 

Bessemer,  343. 

Cementation,  346. 

Composition,  346. 

Crucible,  345. 

Electric,  346. 

Manufacture,  342. 

Open  hearth,  344. 

Properties,  346. 

Special,  346. 

Uses,  347. 
Sterling  silver,  311. 
Stibnite,  283. 
Stove  polish,  156. 
Strontium,  325. 

Compounds,  325. 

Test,  326. 

Sublimation,  272,  301. 
Substances,  i. 
Substitution,  24,  145. 

Valence,  129. 
Sucrose,  194. 
Sugar,  194. 

Beet,  194. 

Cane,  194. 

Fermentation,  196,  203. 


Grape,  196. 

Manufacture,  195. 

Milk,  197. 

Reducing,  197. 

Refining,  195. 

Test,  197. 
Sulphates,  226,  289. 

Acid,  241. 

Normal,  241. 

Test,  242 

Sulphides,  226,  229,  231. 
Sulphites,  234. 

And  sulphur  dioxide,  232. 
Sulphur,  226. 

Amorphous,  228. 

Crystals,  228. 

Dioxide,  232,  390. 

Extraction,  227. 

Flowers,  227. 

Louisiana,  227. 

Modifications,  228. 

Molecular  weight,  123. 

Properties,  227. 

Purification,  227. 

Roll,  227. 

Trioxide,  234. 

Uses,  229. 

Sulphuretted  hydrogen,  229. 
Sulphuric  acid,  235. 

And  copper,  232,  240. 

Chamber  process,  236. 

Contact  process,  238. 

Fuming,  242. 

Manufacture,  235-239. 

Properties,  239. 

Pyro-,  242. 

Test,  142,  242. 

Uses,  241. 

Sulphurous  acid,  233. 
Superphosphate  of  lime,  281. 
Supersaturation,  45. 
Sylvite,  296. 
Symbol,  6. 
Symbols,  60,  inside  back  cover. 

Tallow,  208,  210. 
Tar,  176,  182. 
Tartar,  206. 


INDEX 


409 


Tartar,  cream,  206. 

Emetic,  284. 
Tartaric  acid,  206. 
Tea,  214. 

Temperature,  gas  volumes,  30,  31. 
Normal,  30. 
Standard,  30. 
Tempering,  347. 
Temporary  hardness,  324. 
Tetrad,  127,  129. 
Tetravalent  element,  127. 
Thein,  214. 
Theobromin,  214. 
Theory,  55. 
Atomic,  57. 

Electrolytic  dissociation,  137. 
Thermal,  equation,  161. 
Thermit,  330,  331. 
Thermometer,  389. 

Mercury,  361,  362. 
Thermos  bottle,  197. 
Thomas-Gilchrist  process,  344. 
Thorium,  373,  386. 

Oxide,  190. 
Tin,  365. 

Alloys,  366. 
Block,  38. 
Chlorides,  367. 
Compounds,  366,  367. 
Crystals,  367. 
Disease,  366. 
Metallurgy,  365. 
Oxide,  365,  366,  367. 
Plate,  366. 
Test,  367. 
Travertine,  320. 
Triad,  127,  129. 
Tribasic  acid,  145. 
Trivalent  element,  127. 
Tungsten,  346,  377. 
Turnbull's  blue,  350. 
Tuyeres,  339. 
Type  metal,  284. 

Univalent  element,  127. 
Upward  displacement,  89. 
Uranium,  377. 

And  radium,  385,  386. 


Urea,  217. 

Valence,  126. 

And  atomic  weight,  132. 

Atom,  127. 

Atomic  groups,  127. 

Combination,  129. 

Element,  127,  128. 

Equivalent  weight,  132. 

Exceptions,  131. 

Exercises,  134. 

Formulas,  130. 

Multiple,  130. 

Radicals,  127,  128. 

Representation,  127,  131. 

Substitution,  129. 

Tables,  128. 

Terms,  127. 
Vanadium  steel,  346. 
Vanillin,  214. 
Vapor    density,    molecular    weight, 

"5- 

And  oxygen,  115. 
Vapor  pressure,  39,  40. 

And  deliquescence,  47. 

And  efflorescence,  47. 

Formula,  41. 

Table,  391. 

Temperature,  40. 
Vaseline,  176. 
Venetian  red,  347. 
Vermilion,  363. 
Vinegar,  204,  205. 
Vitriol,  blue,  308. 

Green,  348. 

White,  359. 
Volume,  and  pressure,  30,  32. 

And  temperature,  30,  31. 
Volumetric  equation,  124. 

Water,  36. 

Ammonia,  90. 

And  alcohol,  43. 

And  ions,  147. 

And  sodium,  22,  289. 

Chemical  properties,  41. 

Chlorine,  74. 

Composition,  23. 


4io 


INDEX 


Water,  distilled,  37,  38. 

Drinking,  36,  37. 

Electrolysis,  150. 

Filter,  37. 

Glass,  253. 

Hard,  36,  167,  324,  356. 

In  body,  217. 

Natural,  36. 

Neutralization,  143. 

Occurrence,  36. 

Ocean,  36. 

Of  crystallization,  46. 

Oxygen,  74. 

Physical  properties,  38. 

Purification,  333. 

Rain,  36. 

River,  36. 

Settling,  37,  333. 

Soft,  36,  324. 

Solvent  power,  42. 

Vapor  in  air,  106,  107. 

Vapor  pressure,  39,  391. 
Water  gas,  168,  182. 

Poisonous,  169,  184. 
Weathering,  254. 
Weights  of  gases,  390. 
Welding,  18,  174,  342. 
Welsbach  light,  190. 

Mantle,  190,  373. 
Whey,  198. 
Whisky,  204. 
White  lead,  371. 


Whitewash,  323. 
Wine,  204. 
Wood  alcohol,  202. 
Wood,  petrified,  249. 
Wood's  metal,  285. 
Writing  equations,  67. 
Wrought  iron,  341. 

Xenon,  109. 

Yeast,  196. 

And  alcohol,  203. 
And  bread,  199. 

Zinc,  357. 

Alloys,  359. 

And  nitric  acid,  99. 

Atoms  in  a  molecule,  123. 

Chloride,  359. 

Compounds,  357,  359. 

Granulated,  358. 

Hydroxide,  359. 

Metallurgy,  357. 

Oxide,  358,  359. 

Properties,  358. 

Sulphate,  359. 

Sulphide,  359,  385. 

Tests,  360. 

Uses,  358. 

White,  359. 
Zincates,  358,  360. 
Zymase,  196,  203. 


GENERAL    CHEMISTRY 

PART  II 
EXPERIMENTS 


BY 

LYMAN    C.    NEWELL,   PH.D.    (JOHNS  HOPKINS) 
PROFESSOR  OF  CHEMISTRY  IN  BOSTON  UNIVERSITY 

AUTHOR  OF  "EXPERIMENTAL  CHEMISTRY,"  "DESCRIPTIVE  CHEMISTRY,' 
"INORGANIC  CHEMISTRY  FOR  COLLEGES" 


D.   C.   HEATH  &  CO.,  PUBLISHERS 

BOSTON          NEW  YORK        CHICAGO 


COPYRIGHT,     1914 
BY    LYMAN    C.    NEWELL 


IF4 


PREFACE 

THE  experiments  in  this  book  are  designed  primarily  to 
accompany  the  author's  General  Chemistry  —  Principles  and 
Applications.  They  are  cited  throughout  that  book  and  are 
referred  to  by  number  as  being  in  Part  II.  The  selection, 
arrangement,  and  subdivisions  of  the  experiments  are  such, 
however,  that  teachers  will  find  this  book  itself  serviceable 
under  various  conditions. 

Teachers  should  notice  several  things  about  the  experiments 
in  this  book.  First,  as  a  whole  they  are  divided  into  two 
kinds  —  regular  and  supplementary.  Second,  the  regular 
experiments  include  not  only  those  acknowledged  as  of  funda- 
mental value  in  a  general  course,  but  also  practical  and  novel 
experiments  which  emphasize  the  relation  of  chemistry  to 
everyday  experiences  of  students.  Third,  the  supplementary 
set  includes  experiments  suited  for  beginners,  but  varying 
widely  in  length,  difficulty,  and  utility.  These  experiments  are 
just  what  their  name  implies  —  supplementary.  They  will 
serve  numerous  uses:  e.g.  to  lay  special  emphasis  on  certain 
principles  and  applications,  to  provide  additional  opportunity 
to  acquire  skill  in  manipulation,  to  supplement  information 
gained  by  the  regular  experiments,  to  provide  illustrative 
demonstrations  for  the  classroom,  to  furnish  material  for 
individuals  who  prefer  or  need  special  experiments  or  who  are 
compelled  to  meet  specific  requirements,  to  present  attractive 
laboratory  work  to  students  who  cannot  (or  will  not)  pursue 
chemistry  scientifically,  and  to  stimulate  those  who  are  inter- 
ested in  the  practical  applications  of  chemistry.  Fourth,  this 


iv  PREFACE 

liberal  provision  for  laboratory  work  not  only  permits  the 
selection  of  a  sufficient  number  of  experiments  adapted  to  a 
wide  range  of  equipment,  but  also  enables  teachers  to  accom- 
plish one  or  more  aims,  e.g.  giving  general  mental  training, 
inculcating  the  scientific  point  of  view,  meeting  college  pre- 
paratory requirements,  -teaching  the  fundamental  principles  of 
chemistry,  emphasizing  relations  of  chemistry  to  household 
arts  and  to  industries,  and  utilizing  chemistry  as  a  factor  in 
vocational  education.  These  aims,  varied  as  they  seem,  can 
be  accomplished  by  judicious  utilization  of  text  and  experi- 
ment, not  only  because  adequate  material  is  available  but  also 
because  the  text-book  and  laboratory  manual  have  been  made 
flexible  in  content  and  arrangement.  On  page  xiii  there  are 
suggestions  which  will  assist  teachers  in  selecting  experiments 
suitable  for  different  kinds  of  courses. 

The  directions  for  performing  the  early  experiments  are 
rather  full.  Experience  has  convinced  the  author  that  begin- 
ners need  adequate  directions  at  the  outset.  The  directions 
for  performing  the  semi-quantitative  experiments  are  also  full, 
for  time  and  annoyance  are  saved  by  avoiding  repetitions 
necessitated  by  inadequate  directions. 

The  Introduction  contains  directions  for  preparing,  construct- 
ing, and  arranging  apparatus  and  for  performing  many  general 
laboratory  operations.  Students  should  be  required  to  famil- 
iarize themselves  with  this  part  of  the  book  and  to  use  it  as 
occasion  demands.  In  the  Introduction  there  will  also  be  found 
information  about  the  procedure  in  case  of  accidents  and  some 
suggestions  about  laboratory  note  books.  Annotated  lists  of 
the  supplies  needed  for  the  experiments  will  be  found  at  the 
end  of  the  book. 

L.  C.  N. 

BOSTON,  MASS., 
May,  1914. 


CONTENTS 

(SUPPLEMENTARY  EXPERIMENTS  ARE  MARKED  s.) 

PAGE 

SUGGESTIONS  FOR  TEACHERS xiii 

INTRODUCTION  —  General  Directions  —  Accidents — Note-books  i 

EXPERIMENT 

1.  PROPERTIES  OF  COPPER  SULPHATE n 

2.  PHYSICAL  AND  CHEMICAL  CHANGES 12 

3  s.    PROPERTIES  OF  IODINE 14 

4  s.    PHYSICAL  AND  CHEMICAL  CHANGES 15 

6  s.    PHYSICAL  AND  CHEMICAL  CHANGES 15 

6.  PREPARATION  AND  PROPERTIES  OF  OXYGEN 16 

7.  OXIDATION  OF  COPPER 18 

8  s.    PREPARATION    OF    OXYGEN    FROM    VARIOUS    SUB- 

STANCES       19 

9  s.    HEATING  A  METAL  IN  AIR 20 

10.  PREPARATION  AND  PROPERTIES  OF  HYDROGEN     ...  21 

11.  REDUCTION  OF  COPPER  OXIDE  BY  HYDROGEN     ...  21 

12  s.    PREPARATION    OF    HYDROGEN    FROM   VARIOUS    SUB- 

STANCES       25 

13  s.    WEIGHT  OF  A  LITER  OF  OXYGEN 27 

14.  WATER  IN  FOOD 29 

15.  SOME  PHYSICAL  PROPERTIES  OF  WATER 29 

16.  SOME  CHEMICAL  PROPERTIES  OF  WATER 30 

17.  SOLUBILITY  OF  GASES 32 

18.  SOLUBILITY  OF  LIQUIDS 32 

19.  SOLUBILITY  OF  SOLIDS 33 

20.  CRYSTALLIZATION     34 

21.  TESTING  FOR  WATER  OF  CRYSTALLIZATION 35 

22.  PER  CENT  OF  WATER  OF  CRYSTALLIZATION     ....  35 

23.  EFFLORESCENCE 36 

24.  DELIQUESCENCE 36 


vi  CONTENTS 

EXPERIMENT  PAGE 

25  s.    WATER  IN  SUBSTANCES 37 

26s.    DISTILLED  WATER 37 

27  s.    SOLUBILITY  or  A  SOLID      38 

28  S.      SUPERSATURATION 39 

29.  QUALITATIVE  COMPOSITION  or  WATER 40 

30.  ELECTROLYSIS  OF  WATER      41 

31  s.    HYDROGEN  DIOXIDE 42 

32  s.  COMBINATION  or  OXYGEN  WITH  MAGNESIUM   ....  43 

33.  PREPARATION  AND  PROPERTIES  OF  CHLORINE      ...  45 

34.  BLEACHING  POWDER 47 

35.  PREPARATION    AND    PROPERTIES    OF    HYDROCHLORIC 

ACID 47 

36.  TESTS  FOR  CHLORIDES 49 

37.  GENERAL  PROPERTIES  OF  ACIDS 49 

38.  GENERAL  PROPERTIES  OF  BASES      50 

39.  A  PROPERTY  OF  MANY  SALTS 50 

40.  NEUTRALIZATION 50 

41  s.  PREPARATION    OF    CHLORINE    FROM    VARIOUS    SUB- 

STANCES        51 

42  s.  PREPARATION  OF  HYDROGEN  CHLORIDE  FROM  VARIOUS 

SUBSTANCES 52 

43  s.    AQUA  REGIA     52 

44  s.  LITMUS  REACTION  OF  COMMON  SUBSTANCES     ....  53 

45.  PREPARATION  AND  PROPERTIES  OF  NITROGEN  ....  54 

46.  AMMONIA  AND  AMMONIUM  HYDROXIDE 55 

47.  PREPARATION  OF  NITRIC  Acn>     57 

48.  PROPERTIES  OF  NITRIC  ACID 58 

49.  TEST  FOR  NITRIC  ACID  AND  NITRATES 59 

50.  NITRIC  OXIDE  AND  NITROGEN  DIOXIDE 60 

51.  PROPERTIES  OF  NITRATES 61 

52  s.  PREPARATION    OF    NITROGEN    FROM    VARIOUS    SUB- 

STANCES        62 

53  s.  PREPARATION    OF    AMMONIA    FROM    VARIOUS    SUB- 

STANCES        62 

54  s.  INTERACTION  OF  NITRIC  ACID  AND  METALS     ....  63 

55  s.    SODIUM  NITRITE 63 

56  s.  PREPARATION  AND  PROPERTIES  OF  NITROUS  OXIDE  .    .  64 
57.  PER  CENT  OF  OXYGEN  AND  NITROGEN  IN  AIR    ,  66 


CONTENTS  vii 

EXPERIMENT  PAGE 

58.  WATER  VAPOR  IN  THE  AIR 67 

59.  CARBON  DIOXIDE  IN  THE  AIR 68 

60  s.  TESTING  AIR 68 

61.  EQUIVALENT  OF  ZINC  TO  HYDROGEN 69 

62.  EQUIVALENT  OF  MAGNESIUM  TO  HYDROGEN     ....  71 

63  s.  EQUIVALENT  OF  IRON  TO  COPPER 72 

64  s.  EQUIVALENT  OF  ALUMINIUM  TO  HYDROGEN      «...  72 

65  s.  EQUIVALENT  OF  CALCIUM  TO  HYDROGEN 72 

66.  CHEMICAL  BEHAVIOR  OF  ELECTROLYTES  IN  SOLUTION  73 

67.  CHEMICAL  BEHAVIOR  OF  ELECTROLYTES  IN  SOLUTION  73 

68.  GENERAL  PROPERTIES  OF  ACIDS,  BASES,  AND   SALTS  74 

69.  LITMUS  REACTION  OF  DIFFERENT  SALTS 74 

70.  ELECTROLYSIS  OF  COPPER  SULPHATE 74 

71.  ELECTROLYSIS  OF  SODIUM  SULPHATE      75 

72  s.  ELECTROLYTES  AND  NON-ELECTROLYTES 76 

73  s.  CHEMICAL  BEHAVIOR  OF  ELECTROLYTES  IN  SOLUTION  76 

74  s.  ELECTROLYSIS  OF  POTASSIUM  IODIDE 76 

75  s.  ELECTROLYSIS  OF  WATER 77 

76  s.  COLORED  IONS 77 

77  s.  MIGRATION  OF  IONS 77 

78  s.  NEUTRALIZATION  BY  TITRATION 78 

79  s.  PREPARATION  OF  SALTS 79 

80.  DISTRIBUTION  OF  CARBON 81 

81.  PROPERTIES  OF  COAL      81 

82.  PROPERTIES  OF  CHARCOAL 82 

83.  PREPARATION  AND  PROPERTIES  OF  CARBON  DIOXIDE.  83 

84.  CARBON  DIOXIDE  AND  RESPIRATION 84 

85.  ACID  CALCIUM  CARBONATE .  84 

86.  TESTING  FOR  CARBONATES 85 

87.  PREPARATION  AND  PROPERTIES  OF  ACETYLENE    ...  85 

88  s.  PROPERTIES  OF  GRAPHITE     85 

89  s.  PREPARATION   OF   CARBON   DIOXIDE  ^  BY   DIFFERENT 

METHODS 86 

90  s.  CARBONIC  Aero 87 

91  s.  PREPARATION   AND    PROPERTIES    OF    -CARBON    MON- 

OXIDE        87 

92  s.  PRINCIPLE  OF  THE  DAVY  SAFETY  LAMP 88 

93  s.  PROPERTIES  OF  CARBONUNDUM 89 


viii  CONTENTS 

EXPERIMENT  PAGE 

94.  PREPARATION  AND  PROPERTIES  OF  COAL  GAS  ....  90 

95.  CANDLE  FLAME 90 

96.  BUNSEN  BURNER  AND  FLAME 91 

97.  REDUCTION  AND  OXIDATION  WITH  THE  BLOWPIPE  .   .  93 

98  s.  COMBUSTION  OF  ILLUMINATING  GAS 94 

99  s.  BY-PRODUCTS  OF  ILLUMINATING  GAS  MANUFACTURE  95 

100  s.  TESTING  ILLUMINATING  GAS 95 

101  s.  TESTING  METALS  WITH  THE  BLOWPIPE      95 

102  s.  WELSBACH  BURNER,  MANTLE,  AND  FLAME 95 

103.  COMPOSITION  OF  ORGANIC  COMPOUNDS      96 

104.  PROPERTIES  OF  SUCROSE  (Cane  Sugar)      96 

106.  PROPERTIES  OF  DEXTROSE  (Glucose) 97 

106.  FEHLING'S  TEST  FOR  SUGAR 98 

107.  TESTING  FOR  GLUCOSE 98 

108.  PROPERTIES  OF  STARCH 98 

109.  PROPERTIES  OF  ALCOHOL 99 

110.  PROPERTIES  OF  ACETIC  ACID 100 

111.  TEST  FOR  ACETIC  ACID 100 

112.  PROPERTIES  OF  VINEGAR 100 

113.  TESTING  BAKING  POWDER 100 

114.  GENERAL  PROPERTIES  OF  FATS 102 

116.  PREPARATION  OF  SOAP 102 

116.  PROPERTIES  OF  SOAP 103 

117.  COMPOSITION  OF  PROTEINS 103 

118.  TESTING  FOR  PROTEINS 104 

119.  TESTING  FOOD 105 

120.  TESTING  FLOUR 105 

121  s.  PREPARATION  OF  INVERT  SUGAR  FROM  SUCROSE     .   .  106 

122  s.  TESTING  FOR  SUGAR  IN  VEGETABLES  AND  FRUITS  .   .  106 

123  s.  DETECTION  OF  STARCH  BY  IODINE 106 

124  s.  CONVERSION  OF  STARCH  INTO  SUGAR  BY  AN  ENZYME  107 
126  s.  PROPERTIES  OF  DEXTRIN 107 

126  s.  GUNCOTTON,  COLLODION,  AND  CELLULOID 107 

127  s.  PREPARATION  AND  PROPERTIES  OF  ETHYL  ALCOHOL  108 

128  s.  FORMALDEHYDE 109 

129  s.  ETHER 109 

130.  PHYSICAL  PROPERTIES  OF  SULPHUR no 

131.  PREPARATION  OF  CRYSTALLIZED  SULPHUR  no 


CONTENTS  ix 

EXPERIMENT  PAGE 

132.  PREPARATION  or  AMORPHOUS  SULPHUR     in 

133.  CHEMICAL  PROPERTIES  OF  SULPHUR in 

134.  SULPHUR  DIOXIDE  AND  SULPHUROUS  ACID 112 

135.  PROPERTIES  OF  SULPHURIC  Acn>     113 

136.  TESTS  FOR  SULPHURIC  ACID  AND  SULPHATES   ....  114 

137s.    SULPHUR  MATCHES     115 

138  s.    PREPARATION  AND  PROPERTIES  OF  HYDROGEN  SUL- 
PHIDE    115 

139s.    PREPARATION  AND  PROPERTIES  OF  SULPHIDES      ...  116 

140s.    PROPERTIES  OF  SULPHUROUS  Aero 117 

141  s.    TESTS  FOR  SULPHUR 117 

142.  CRYSTALLIZATION  OF  BORAX 118 

143.  PROPERTIES  OF  BORAX 118 

144.  TESTS  WITH  BORAX  BEADS 118 

145  s.    PREPARATION  AND  PROPERTIES  OF  BORIC  ACID   ...  120 

146.  PROPERTIES  OF  SILICON 121 

147.  TEST  FOR  SILICON 121 

148s.    PREPARATION  AND  PROPERTIES  OF  SILICIC  ACID    ...  121 

149  s.    THE  CYCLE  OF  SILICON  DIOXIDE 122 

150  s.    TESTING  FOR  SILICON 123 

151  s.    PROPERTIES  OF  GLASS 123 

162.       PREPARATION     AND     PROPERTIES     OF      HYDROGEN 

FLUORIDE 124 

153.  PREPARATION  AND  PROPERTIES  OF  BROMINE    ....  125 

154.  PREPARATION     AND     PROPERTIES     OF     MAGNESIUM 

BROMIDE 126 

155.  TESTS  FOR  BROMINE 126 

156.  PREPARATION  AND  PROPERTIES  OF  IODINE 127 

157.  TESTS  FOR  FREE  IODINE 128 

158.  TESTS  FOR  IODINE  IN  IODIDES 128 

159.  TESTS    FOR    ORTHOPHOSPHORIC    ACID    AND    ORTHO- 

PHOSPHATES       129 

160.  TESTS    FOR    METAPHOSPHORIC    ACID    AND    META- 

PHOSPHATES       129 

161.  ARSENIC  TRISULPHIDE 130 

162.  ANTIMONY  TRICHLORIDE 131 

163.  ANTIMONY  TRISULPHIDE , 131 

164s.    PROPERTIES  OF  PHOSPHORUS    . 131 


x  CONTENTS 

EXPERIMENT  PAGE 

165  s.  PROPERTIES  OF  ANTIMONY 132 

166  s.  INTERACTION  or  ANTIMONY  AND  ACIDS      132 

167  s.  PROPERTIES  OF  BISMUTH 133 

168  s.  BISMUTH  TRICHLORIDE 133 

169  s.  FUSIBLE  ALLOYS 133 

170.  TESTS  FOR  SODIUM 134 

171.  PROPERTIES  OF  SODIUM  CHLORIDE 134 

172.  PROPERTIES  OF  SODIUM  HYDROXIDE    .    . 135 

173.  PROPERTIES  OF  POTASSIUM 135 

174.  TESTS  FOR  POTASSIUM 136 

175.  PROPERTIES  OF  AMMONIUM  CHLORIDE 136 

176  s.  PROPERTIES  OF  SODIUM 136 

177  s.  SODIUM  BICARBONATE 137 

178  s.  TESTING  FOR  SODIUM  AND  POTASSIUM  CARBONATES   .  138 

179  s.  POTASSIUM  NITRATE 138 

180  s.  PROPERTIES  OF  AMMONIUM  COMPOUNDS 138 

181.  PROPERTIES  OF  COPPER      139 

182.  TESTS  FOR  COPPER     139 

183.  PROPERTIES  OF  COPPER  SULPHATE 140 

184.  DISPLACEMENT  OF  METALS  —  COPPER 140 

185.  TESTS  FOR  SILVER 141 

186.  PROPERTIES  OF  GOLD      141 

187.  TEST  FOR  GOLD 142 

188  s.  TESTS  FOR  COPPER  IN  ALLOYS 142 

189  s.  CUPROUS  OXIDE 143 

190  s.  DEPOSITION  OF  A  SILVER  FILM -  .    .    .    .  143 

191  s.  DISPLACEMENT  OF  METALS  —  SILVER 143 

192  s.  TARNISHING  AND  CLEANING  SILVER 143 

193  s.  SILVER  HALIDES 143 

194  s.  TESTING  FOR  COPPER,  SILVER,  AND  GOLD 144 

195.  PROPERTIES  OF  CALCIUM ' 145 

196.  TESTS  FOR  CALCIUM 146 

197.  TESTING  FOR  CALCIUM 146 

198.  PLASTER  OF  PARIS      146 

199.  HARD  WATER 147 

200.  TESTS  FOR  STRONTIUM 147 

201.  TESTS  FOR  BARIUM 147 

202  s.  ACID  CALCIUM  CARBONATE    .                                   .   .  148 


CONTENTS  xi 

EXPERIMENT  PAGE 

203  s.    CALCIUM  OXIDE  AND  HYDROXIDE 148 

204  s.    RED  AND  GREEN  FIRE '149 

205.  PROPERTIES  or  ALUMINIUM 150 

206.  ALUMINIUM  HYDROXIDE 1 50 

207.  CLARIFICATION  OF  WATER 151 

208.  THERMIT 151 

209.  TESTS  FOR  ALUMINIUM 152 

210  s.    ALUMINIUM  SALTS  AS  MORDANTS 152 

211  s.    POTASSIUM  ALUM 152 

212s.    DISPLACEMENT  OF  METALS  —  ALUMINIUM 153 

213s.    EQUIVALENT  OF  ALUMINIUM      153 

214  s.    HYDROLYSIS  OF  ALUMINIUM  SALTS 153 

215  s.    ALUM  BAKING  POWDER 153 

216.  PROPERTIES  OF  IRON 154 

217.  FERROUS  COMPOUNDS     154 

218.  FERRIC  COMPOUNDS 155 

219.  INTERRELATION  OF  IRON  COMPOUNDS 155 

220  s.    TESTING  FOR  IRON 156 

221  s.    BLUE  PRINT  PAPER 156 

222  s.    TEST  FOR  NICKEL 157 

223  s.    TEST  FOR  CCBALT 157 

224.  TESTS  FOR  MAGNESIUM     158 

225.  TESTS  FOR  ZINC ,  •    •    •  158 

226.  MERCUROUS  AND  MERCURIC  COMPOUNDS 159 

227  s.    PROPERTIES  OF  MAGNESIUM  AND  ZINC 159 

228  s.    EQUIVALENT  OF  MAGNESIUM  AND  ZINC     159 

229  s.    PHYSICAL  PROPERTIES  OF  MERCURY 160 

230  s.  DISPLACEMENT  —  MAGNESIUM,  ZINC,  AND  MERCURY  .  160 

231  s.    TEST  FOR  CADMIUM    ., 160 

232.  TEST  FOR  TIN 161 

233.  TESTS  FOR  LEAD      161 

234  s.    PROPERTIES  OF  TIN  AND  LEAD 161 

235s.    DISPLACEMENT  —  TIN  AND  LEAD 162 

236  s.    TESTING  FOR  TIN  AND  LEAD 162 

237  s.    ANALYSIS  OF  SOLDER .    .  162 

238  s.  SEPARATION  OF  LEAD,  SILVER,  AND  -MERCURY    ...  163 

239.  TESTS  FOR  CHROMIUM. 164 

240.  POTASSIUM  CHROMATE  AND  DICHROMATE      164 


xii  CONTENTS 

EXPERIMENT  PAGE 

241.       TESTS  FOR  MANGANESE 165 

242  s*    CHROMATES  AND  CHROMIC  COMPOUNDS 165 

243  s.     CHROMIC  HYDROXIDE 166 

244  s.    OXIDATION  WITH  POTASSIUM  PERMANGANATE   ....  166 

LABORATORY  EQUIPMENT     167 


SUGGESTIONS   FOR   TEACHERS 

The  author  does  not  intend  that  all  the  experiments  in  this  book 
shall  be  performed  by  a  class  in  the  time  usually  devoted  to  a  first 
course  in  chemistry.  In  selecting  experiments  well  adapted  to 
equipment  and  best  suited  to  needs,  teachers  will  find  the  subjoined 
lists  serviceable.  The  numbers  in  parentheses  refer  to  optional 
experiments. 

LIST  I  —  General  Course.  This  group  contains  experiments 
calling  for  a  well  equipped  laboratory  and  a  liberal  amount  of  time, 
and  as  far  as  possible  the  course  should  include  these  experiments. 
Numbers  1,  2,  (4s),  (5s),  6,  7,  (8s),  (9s),  10,  11,  (12s),  (14),  (16),  16, 
17,  18,  19,  20,  21,  23,  24,  26s,  (28s),  (29),  33  or  41s,  34,  35  or  42s,  36, 
37,  38,  39,  40,  (43s),  (44s),  45  or  52s,  46  or  53s,  (47),  (48),  49,  50, 
(54s),  (56s),  58,  59,  66  or  67,  70  or  71,  79s,  80,  81,  82,  83,  84,  85,  (88s), 
(89s),  (91s),  (94),  (95),  (96),  97,  (104),  (106),  (109),  (111),  (115), 
(116),  131,  132,  (133),  134,  (135),  136,  (138s),  139s,  (140s),  144,  153, 
156,  (157),  (158),  (159),  170,  (171),  174,  (176s),  177s,  179s,  181,  182, 
184,  185,  (187),  (191s),  (193s),  196,  (197),  (198),  199;  (200),  (201), 
203s,  (204s),  (206),  (207),  209,  (212s),  217,  218,  219,  (225),  (229s), 
(230s),  (235s),  238s. 

LIST  II  —  Shorter  Course.  The  experiments  in  this  group  con- 
stitute a  consecutive  course  and  are  suitable  wherever  equipment  is 
moderate  and  time  limited.  Numbers  1,  2,  6,  10,  14,  17,  18,  19,  (20), 
23,  24,  34,  35  or  42s,  36,  37,  38,  39,  40,  (41s),  46  or  53s,  49,  50,  79s, 
80,  83  or  89s,  94,  95,  (96),  (97),  104,  106,  108,  (111),  (112),  (115), 
(116),  (123s),  131,  132,  135,  136,  (144),  (153),  (156),  (157),  (158), 
170,  174,  179s,  182,  184,  196,  199,  203s,  (204s),  207,  209,  217,  218, 
219,  (220). 

LIST  III  —  College  Preparatory  Course.  This  group  covers  the 
usual  entrance  requirement  for  college.  Teachers  are  urged  to 
substitute  short  experiments  from  the  supplementary  set  wherever 
the  principle  is  identical.  Numbers  (1),  (2  or  4s  or  5s),  6,  7,  8s,  9s, 
10,  11,  12s,  13s,  (14  or  25s),  (15),  16,  (17),  (18),  (19),  (20),  21,  22, 
(23),  (24),  26s,  29,  (32s),  33,  35,  36,  40,  (41s),  (42s),  46,  47,  48,  49, 


xiv  CHEMISTRY 

50,  (52s),  (53s),  (54s),  56,  57,  61  or  64s,  62  or  65s,  (63s),  (66),  (67  or 
73s),  (68),  (70  or  71),  78s,  79s,  80,  81,  82,  83,  84,  85,  88s,  (89s),  90s, 
91s,  (94s),  (95s),  (96s),  (97s),  99s,  (104),  (105),  (106),  (108),  (109), 
111,  (1^2),  115,  116,  (117),  (118),  119,  120,  127s,  131,  132,  134,  (135), 
136,  138s,  139s,  (144),  (152),  153,  (165),  156,  (157),  (158),  (159), 
(160),  (170),  (171),  172,  176s,  177s,  (178s),  179s,  (181),  182,  184, 
(185),  (186),  (191s),  (193s),  196,  (197),  (198),  199,  200,  201,  203s, 
(206),  (207),  209,  (212s),  217,  218,  219,  220s,  (223s),  (225),  (229s), 
(230s),  (235s),  238s. 

LIST  IV  —  Practical  Course.  This  group  includes  experiments 
which  emphasize  the  applications  of  chemistry.  The  fundamental 
experiments  in  List  II  may  be  substituted  for  certain  ones  in  this 
list;  omissions  may  also  be  made  as  time  and  equipment  determine. 
Numbers  14,  15,  17,  18,  19,  (25s),  26s,  31s,  34,  36,  37,  38,  39,  44s, 
49,  52s,  53s,  58,  59,  60s,  69,  70,  79s,  80,  81,  82,  83,  86,  87,  88s,  (89s), 
92s,  93s,  94,  98s,  99s,  100s,  101s,  102s,  106,  107,  111,  112,  113,  115, 
116,  118,  119,  120,  122s,  123s,  124s,  126s,  128s,  136,  137,  140s,  150s, 
151s,  159,  169s,  171,  172,  175,  177s,  178s,  179s,  180s,  188s,  190s, 
192s,  194s,  197,  198,  199,  203s,  204s,  207,  208,  210s,  216,  220s,  221s, 
225,  232,  233,  236s,  237s,  238s. 

LIST  V  —  Food  Experiments.  The  fundamental  experiments  on 
the  chemistry  of  food  are  collected  in  this  group.  Numbers  14,  25s, 
26s,  84,  103,  104,  105,  106,  107,  108,  109,  110,  111,  112,  113,  114, 
115,  116,  117,  118,  119,  120,  121s,  122s,  123s,  124s,  125s,  127s. 

LIST  VI  —  Quantitative  Experiments.  This  group  includes  the 
experiments  that  involve  accurate  weighing  and  measuring.  Num- 
bers 13s,  22,  27s,  32s,  57,  61,  62,  63s,  64s,  65s,  78s,  88s(7>),  130(6), 
136(o),  146(6),  151s(d),  165s(6),  181(c),  186(c),  213s,  216(c),  228s, 
229(c). 

LIST  VII  —  Demonstration  Experiments.  In  this  group  are  placed 
experiments  which  may  be  performed  by  he  teacher  before  the  class. 
Many  of  these  experiments  may  be  used  to  supplement  the  indi- 
vidual work  suggested  in  List  II  or  as  substitutes  for  certain  experi- 
ments in  List  I.  Numbers  4s,  5s,  14,  26s,  28s,  29,  30,  34,  43s,  44s, 
47,  48,  50,  66,  67,  68,  70,  71,  72s,  73s,  74s,  76s,  77s,  85,  90s,  92s,  95, 
98s,  102s,  127s,  132,  133,  140s,  142,  145s,  148s,  154,  177s,  190s, 
192s,  198,  202s,  204s,  207,  208,  210s,  239,  240,  241,  244s. 


EXPERIMENTS 


GENERAL   CHEMISTRY 

EXPERIMENTS 

INTRODUCTION 

GENERAL    DIRECTIONS  —  ACCIDENTS  —  NOTE- 
BOOKS 

1.  The  Bunsen  burner  is  used  as  the  source  of  heat  in  most 
chemical  laboratories.     It  is  attached  to  the  gas  cock  by  a 
piece  of  rubber  tubing.     It  is  lighted  by  turning  on  the  gas 
full  and  then  holding  a  lighted  match  in  the  gas  a  short  dis- 
tance above  the  top  of  the  burner.     If  the  flame  is  yellow, 
turn  the  ring  at  the  bottom  of  the  burner  until  the  flame  is  a 
faint  blue.     The  colorless  or  bluish  flame  should  be  used  in 
all  experiments  unless  the  directions  state  otherwise.     The 
hottest  part  of  the  flame  is  near  the  top. 

2.  Heating.  —  The  following  directions  should  be  observed 
in  heating  with  the  Bunsen  burner:  — 

(1)  The  burner  should  always  be  lighted  before  any  piece 
of  apparatus  is  held  over  it,  or  before*  it  is  placed  beneath  a 
wire  gauze  which  supports  a  dish  or  flask. 

(2)  Glass  and  porcelain  apparatus  should  not  be  heated 
when  empty  nor  over  a  bare  or  free  flame  even  if  they  contain 
something  —  unless  directions  so  state.     Vessels  requiring  a 
support   should  be  placed  on  a  wire  gauze  which  stands  on 
the  ring  of  an  iron  stand,  and  heated  gradually  from  beneath 
(Fig.  no).     Vessels  should  be  heated  and  cooled  gradually; 
if  removed  from  the  gauze  while  hot,  they  should  be  placed  on 
a  block  of  wood  or  piece  of  asbestos  —  never  on  a  cold  surface. 

(3)  Many  experiments  require  the  heating  of  test  tubes. 
These  tubes  should  be  dry  on  the  outside.     The  temperature 


CHEMISTRY 


of  a  test  tube  containing  a  solid  should  be  raised  gradually 
by  moving  it  in  and  out  of  the  flame,  or  by  holding  it  in  the 
flame  and  rolling  it  slightly  between  the  thumb  and  forefinger. 
Special  care  must  be  taken  to  distribute  the  heat  evenly.  If 
the  test  tube  contains  a  liquid,  as  is  usually  the  case,  only 
that  part  containing  the  liquid  should  be  heated;  the  test 
tube  should  also  be  inclined  so  that  the  greatest  heat  is  not 
applied  to  the  thin  bottom.  When  the  liquid  begins  to  boil, 
the  test  tube  should  be  removed  from  the  flame  for  an  instant 
or  held  over  it.  In  some  experiments  test  tubes  can  be  held 
between  the  thumb  and  forefinger  without  discomfort.  As  a 
rule  a  test  tube  holder  should  be  used  (Fig.  100). 

3.  Cutting,  Bending,  and  Drawing  Glass  Tubing.  —  (a) 
Cutting.  Determine  the  length  needed,  lay  the  tube  on  the 
desk,  and  with  forward  strokes  of  a  triangular  file  make  a 
short  but  deep  scratch  where  the  tube  is  to  be  cut.  Grasp  the 
tube  in  both  hands,  and  hold  the  thumbs  together  behind  the 
scratch;  now  push  gently  with  the  thumbs,  pull  at  the  same 
time  with  the  hands,  and  the  tube 
will  break  at  the  desired  point.  The 
sharp  ends  should  be  smoothed  by  rub- 
bing them  with  emery  paper  or  by 
rotating  them  slowly  in  the  Bunsen 
flame  until  a  yellow  color  is  distinctly 
seen  or  until  the  end  becomes  red- 
hot. 

(b)  Bending.  Glass  tubes  are  bent 
in  a  flat  flame.  An  ordinary  illumi- 
nating gas  flame  may  be  used,  but  the 
Bunsen  flame  can  be  flattened  by  a 
wing-top  attachment,  which  slips  over  the  top  of  the  burner 
tube.  The  flattened  Bunsen  flame  should  be  slightly  yel- 
low and  about  7  centimeters  (2.5  inches)  wide  for  ordi- 
nary bends.  A  right-angle  bend  is  made  as  follows: 
Determine  the  point  at  which  the  tube  is  to  be  bent. 
Grasp  the  tube  in  both  hands,  and  hold  it  so  that  the  part  to 


Fig.  ioo. — Test  Tube 
and  Holder. 


GENERAL   DIRECTIONS  3 

be  bent  is  directly  over  the  flame.  Slowly  rotate  it  between 
the  thumbs  and  forefingers,  and  gradually  lower  it  into  the 
flame.  Continue  to  rotate  it  until  the  glass  feels  soft  and 
ready  to  yield.  Then  remove  it  from  the  flame,  and  slowly 
bend  it  into  a  right  angle.  It  is  convenient  to  have  at  hand 
a  block  of  wood  or  some  other  right-angled  object  to  assist 
the  eye  in  completing  the  bend  into  an  exact  right  angle.  It 
is  desirable  though  not  always  necessary  to  anneal  the  bent 
part  of  the  tube.  This  is  done  by  holding  it  in  a  yellow  flame 
until  it  becomes  coated  with  soot;  it  should  then  be  placed  on 
a  block  of  wood,  and  when  cold  wiped  clean.  Tubes  can  be 
bent  into  an  oblique  angle  by  heating  them  through  about 
twice  the  space  required  for  a  right  angle;  a  very  slight  bend, 
however,  is  often  made  by  holding  the  tube  across  the  flame 
and  heating  a  short  space. 

(c)  Drawing.  Glass  tubes  can  be  drawn  to  a  finer  bore  or 
into  two  pointed  tubes  as  follows:  Heat  the  tube  as  in  (b) 
through  about  2.5  centimeters  (i  inch)  of  its  length,  remove 
from  the  flame  and  slowly  pull  it  apart  a  short  distance;  let 
it  cool  for  a  few  seconds,  and  then  pull  it  quickly  to  the  desired 
length.  Stirring  rods  can  be  made  from  glass  rod  in  the  same 
wayt. 

4.  Filtering.  —  A  solid  may  be  separated  from  a  liquid  by 
filtering.  A  circular  piece  of  porous  paper  is  folded  to  fit  a 
glass  funnel,  and  when  the  mixture  is  poured  upon  this  paper 
the  solid  —  the  residue  or  precipitate  —  is  retained,  while 
the  liquid  —  the  filtrate  —  passes  through  and  may  be  caught 
in  a  test  tube  or  any  other  vessel.  The  filter  paper  is  prepared 
for  the  funnel  by 
folding  it  success- 
ively into  the  shapes 

shown  in  Fig.  101 —  x  n 

I,  II,  and  then  open-  Fig.  101.  —  Folded  Filter  Paper, 

ing  the  folded  paper  so  that  three  thicknesses  are  on  one  side 
and  one  on  the  other  as  in  III.  The  cone-shaped  paper  is 
next  placed  in  the  funnel  and  moistened  with  water,  so 


CHEMISTRY 


that  it  will  stick  to  the  sides  of  the  funnel.    The  liquid  to  be 
filtered  may  be  poured  directly  upon  the  paper  or  down  a 

glass  rod  which 
touches  the  edge  of 
the  test  tube;  the 
lower  end  of  the  rod 
should  nearly  touch 
the  paper  inside  the 
funnel,  so  the  liquid 
will  run  down  the 


side    and     thereby 
avoid  bursting  the 
The  funnel  can  be  supported  as 


Fig.  102.  —  Funnel  Supported  for  Filtering. 

apex  of  the  filter  paper, 
shown  in  Fig.  102. 

5.  Constructing  and  Arranging  Apparatus.  —  The  various 
parts  of  an  apparatus  should  be  collected,  prepared,  and 
put  together  before  starting  the  experiment  in  which  the  ap- 
paratus as  a  whole  is  used.  The  parts  that  are  to  fit  each 
other  should  be  selected  and  arranged  so  that  all  joints  are 
gas-tight,  and  as  a  final  precaution,  especially  in  long  experi- 
ments or  those  involving  weighing,  the  apparatus  should  be 
approved  by  the  Teacher.  The  following  suggestions  will  be 
helpful: - 

(1)  To  insert  a  glass  tube  into  rubber  tubing.     Cut  one  end 
of  the  rubber  tubing  at  an  angle,  moisten  the  smoothed  end 
of  the  glass  tube  with  water,  place  the  end  of  the  glass  tube 
in  the  angular-shaped  cavity  so  that  both  tubes  are  at  about 
a  right  angle,  grasp  the  rubber  tube  firmly  and  slip  it  slowly 
up  and  over  the  end  of  the  glass  tube. 

(2)  To  fit  a  glass  tube  to  a  stopper.     Moisten  one  end  of  the 
tube  with  water  and  grasp  it  firmly  near  this  end;   hold  the 
stopper  between  the  thumb  and  forefinger  of  the  other  hand, 
and  work  the  tube  into  the  hole  by  a  gradual  rotary  motion. 
Proceed  in  the  same  manner,  if  the  tube  is  to  be  pushed 
through  the  stopper.      Never  point  the  tube  toward  the  palm 
of  the  hand  that  holds  the  stopper.     Never  grasp  a  bent  tube 


GENERAL   DIRECTIONS  $ 

at  the  bend  when  inserting  it  into  a  stopper  —  it  may  break 
and  cut  the  hand  severely. 

(3)  To  bore  a  hole  in  a  cork.     Rubber  stoppers  are  preferable, 
but  if  corks  are  used,  they  can  be  bored  as  follows:   Select  a 
cork  free  from  cracks  or  channels  and  use  a  borer  which  is 
one  size  smaller  than  the  desired  hole.     Hold  the  cork  between 
the  thumb  and  forefinger,  press  the  larger  end  against  a  firm 
but  soft  board,  and  slowly  push  the  borer  (previously  mois- 
tened with  water  or  soap  solution)  by  a  rotary  movement 
through  the  cork,  taking  care  to  bore  perpendicularly  to  the 
cork.      If  the  hole  is  too  small,  enlarge  it  with  a  round  file. 

(4)  To  make  a  test  wire,     (a)  Platinum.     Rotate  one  end  of 
a  piece  of  glass  rod,  about  10  centimeters  (4  inches)  long,  in 
the  flame  until  it  softens.     At  the  same  time  grasp  a  pieqe  of 
platinum  wire  about  7  centimeters  (3  inches)  long  firmly  in 
the  forceps  about  i  centimeter  (.5  inch)  from  the  end,  and 
hold  it  in  the  flame.     When  the  rod  is  soft  enough,  gently 
push  the  hot  wire  into  the  rod.     If  a  glass  tube  is  used,  it 
should  be  drawn  out  to  a  very  small  diameter  (see  3  (c)) 
before   inserting   the   platinum  wire,  but   in   other   respects 
the  two  operations  are  practically  identical,     (b)    Nichrome. 
An  efficient  test  wire  for  many  experiments  can  be  made  by 
winding   a   piece   of   nichrome  wire  around  a  match  stick. 
The  completed  wires  are  shown  in  Fig.  103. 


Fig.  103.  —  Test  Wires —  Platinum  (Upper),  Nichrome  (Lower). 

6.  Manipulation.  —  Ability  to  use  apparatus  rapidly,  ac- 
curately, and  neatly  is  acquired  only  by  experience.  The 
following  suggestions  will  facilitate  the  acquisition  of  this 
needful  skill:  — 

(i)  Pouring  liquids  and  transferring  solids,  (a)  Liquids 
can  be  poured  from  a  vessel  without  spilling,  by  moistening 


CHEMISTRY 


a  glass  rod  with  the  liquid  and  then  pouring  it  down  the  rod. 

The  angle  at  which  the  rod  is  held  varies  with  circumstances. 

This  is  a  convenient  way  to  pour  a  liquid  from  a  vessel  con- 
taining a  solid  with- 
out disturbing  the 
solid.  (b)  Liquids 
should  be  poured  from 
a  .bottle  by  holding 
the  bottle  as  shown  in 
Fig.  104.  Notice  that 
the  stopper  and  bottle 
are  held  in  the  same 

Fig.  104.  -  Pouring  a  Liquid  from  a  Bottle.     hand-     The  stopper  is 

removed   by   holding 

palm  of  the  hand  upward  and  grasping  the  stopper  between  the 
the  fingers  before  the  bottle  is  lifted  (Fig.  105).  All  stoppers 
should  be  removed  this  way  when  possible,  and  not  laid  down, 
because  the  impurities  adhering  to  the  stopper  may  run 
down  into  the  bottle  and  contaminate  the  solution.  The 
drop  on  the  lip  of  the  bottle  should  be  touched  with  the 
stopper  before  the  latter  is  put  into  the  bottle;  this  simple 
operation  prevents  the  drop  from 
running  down  the  outside  of  the 
bottle  upon  the  label  or  upon  the 
shelf,  (c)  Solids  should  never  be 
poured  directly  from  a  large  bottle 
into  a  test  tube,  retort,  or  similar 
vessel.  A  convenient  method  is  as 
follows:  Rotate  the  bottle  slowly 
so  that  the  solid  will  roll  out  in 
small  quantities;  catch  the  solid 
on  a  narrow  strip  of  paper  creased 


Fig.  105.  —  Removing  the 
Stopper  from  a  Bottle. 


lengthwise,  and  slide  the  solid  from  the  paper  into  the  desired 
vessel. 

(2)    Collecting    gases.     Gases    are    usually    collected    over 
water  by  means  of  a  pneumatic  trough,  a  common  form  of 


GENERAL   DIRECTIONS  7 

which  is  shown  in  Fig.  108.  The  vessel  to  be  filled  with  gas 
is  first  filled  with  water,  covered  with  a  piece  of  filter  paper, 
inverted,  and  placed  mouth  downward  on  the  support  of  the 
trough,  which  is  previously  filled  with  water  just  above  the 
support.  The  paper  is  then  removed,  and  the  vessel  slipped 
over  the  hole  in  the  support.  Glass  plates  instead  of  filter 
paper  may  be  used  to  cover  the  bottle.  The  gas  which  is 
evolved  in  the  generator  passes  through  the  delivery  tube, 
and  bubbles  up  through  the  water  into  the  vessel,  forcing  the 
water  out  of  the  vessel  as  it  rises.  All  gases  insoluble  in  water 
may  be  thus  collected.  Some  heavy  gases,  such  as  hydro- 
chloric acid,  chlorine,  and  sulphur  dioxide,  are  collected  by 
allowing  the  gas  to  flow  downward  into  an  empty  bottle,  and 
displace  the  air  in  the  bottle,  i.e.  by  downward  displacement 
(Fig.  122).  Ammonia  and  other  light  gases  are  usually  col- 
lected by  allowing  the  gas  to  flow  upward  into  a  bottle,  i.e. 
by  upward  displacement  (Fig.  125). 

7.  Weighing.  —  Most  experiments  in  this  book  involve 
only  approximate  weights  of  substances;  a  few  require  ac- 
curate weights.  Approximate  weighings  are  made  on  the 
scales  and  accurate  weighings  on  the  balance. 

The  following  rules  should  be  observed  in  all  weighings:  — 

(a)  Before  weighing,  see  that  the  scales  and  balance  are 
clean  and  properly  adjusted.     If  out  of  order,  do  not  attempt 
the  adjustment  yourself,  but  report  the  case  to  the  Teacher. 

(b)  Substances  are  put  on  the  left  side  and  weights  on  the 
right.     Heavy  objects  and  weights  should  be  put  in  the  center 
of  the  pan. 

(c)  Substances  should  not  be  placed  directly  on  the  platform 
or  pan,  except  pieces  of  metal  or  glass  objects.     In  weighing 
on  the  scales,  put  pieces  of  paper  of  about  the  same  size  in 
each  platform;    the  left  one  should  be  creased.     Take  the 
substance  from  the  bottle  with  a  clean  spoon  or  spatula,  or 
pour  it  out  by  rotating  the  bottle  as  described  in  6  (c);   if 
you  weigh  out  too  much,  do  not  put  it  back  into  the  bottle, 
but  throw  it  into  the  waste  jar  or  a  special  bottle.     In  using 


8  CHEMISTRY 

the  balance,  if  the  substance  should  not  be  placed  on  the  pan, 
weigh  a  small  watch  crystal  or  crucible  and  then  weigh  the 
substance  in  this  vessel.     Sometimes  a  piece  of  apparatus  is 
not  put  on  the  pan  but  hung  from  the  balance  hook. 
The  process  of  weighing  is  as  follows :  — 

A.  Scales.     Put  the  object  or  the  paper  and  substance  on 
the  left  side;    on  the  right  side  put  the   exact  weight  if  it  is 
known  or  the  approximate  weight  if  the  exact  weight  is  not 
known.     Now  add  or  remove  substance  or  weights  until  the 
pointer  swings  the  same  number  of  spaces  each   side  of  the 
middle  division.      Weighings  of  single  grams  and  fractions  are 
usually  made  by  sliding  a  rider  along  a  graduated  beam  on 
the  front  of  the  scales. 

B.  Balance.     Put  the  substance  or  object  on  the  left  pan 
and  the  weight  judged  to  be  equal  on  the  right  pan.     Release 
the  beam  carefully  by  turning  the  screw  or  lever,  and  note 
the  movement  of  the  pointer.     If  the  added  weight  is  incorrect, 
arrest  the  beam  and  change  the  weights,  taking  care  to  add 
or  remove  the  weights  systematically.     Then  release  the  beam 
again  and  observe  as  before;  if  the  pointer  does  not  swing  the 
same  number  of  spaces  each  side  of  the  central  line,  arrest 
the   beam   and   change   the   weights   accordingly.     Continue 
until  the  correct  weight  is  obtained'.     As  soon  as  the  substance 
or  object  is  weighed,  note  the  weights  on  the  pan  and  record 
at  once,  then  compare  the  weights  with  those  missing  from 
the  box;   if   correct,  so  indicate  in  the  notebook,  and   finally 
check  the  weight  by  noting  the  weights  as  they  are  returned 
to  the  box.     The  following  should  be  rigidly  observed :  - 

(a)  Always  arrest  the  beam  before  changing  the  weights  or 
the  load  (i.e.  the  object  or  substance). 

(b)  If  on  releasing,  the  beam  does  not  swing,  arrest  and 
release  again,  or  fan  one  pan  very  gently. 

(c)  Record  the  result  of  all  weighings  in  the  proper  place 
in  a  notebook,  —  never  on  a  scrap  of  paper. 

(d)  Handle  all  weights  with  special  forceps,  unless  otherwise 
directed. 


GENERAL   DIRECTIONS  9 

8.  Measuring.  —  Liquids  are  measured  in  graduated  cylin- 
ders and  burettes  (Figs.  134,  132).     The  lowest  point  of  the 
curved  surface  of  the  liquid,  called  the  meniscus,  is  its  correct 
height   (Fig.   106).     The  average  or- 
dinary test  tube    (6  X  f  inch)  holds 

about    30    cubic    centimeters,    while 

the  large  test  tube  (8  X  i  inch)  holds  ^ 

about    75   cubic    centimeters.     Time     ^  __ 

can  be  saved  by  remembering  these 

volumes. 

9.  Smelling    and    Tasting.  —  Un- 
familiar  substances    should  never  be 

tasted  or  smelled  except  according  to  Fig.  106.  —  Meniscus.  Cor- 
directions,  and  even  then  with  the  rect  ReadinS  is  along 
utmost  caution.  Never  inhale  a  gas 

vigorously,  but  waft  it  gently  with  the  hand  toward  the  nose. 
Taste  acids,  etc.,  by  touching  a  minute  portion  of  the  dilute 
solution  to  the  tip  of  the  tongue,  and  as  soon  as  the  sensation 
is  detected,  reject  the  solution  at  once  —  never  swallow  it. 

10.  Accidents.  —  (i)  Cuts  should  be  washed  in  clean  cold 
water  and  then  covered  with  collodion  or  court  plaster  if 
slight,  or  bandaged  firmly  if  severe.     (2)  Burns  caused  by 
hot  objects  should  be  covered  with  a  paste  made  by  mixing 
sodium  bicarbonate  (baking  soda)  and  carron  oil  (an  emulsion 
of  lime  water  and  oil)  and  then  bandaged.     (3)  Acids  and 
alkalies  if  spilled  on  the  hands  or  spattered  on  the  face  should 
be  washed  off  with  water;   if  a  burn  is  produced,  this  may  be 
treated  as  described  above.     (4)  If  a  poison  is  swallowed,  a 
physician  should  be  called  at  once;    meanwhile  an  emetic 
consisting  of  warm  water  and  mustard  should  be  administered, 
and  subsequently  the  proper  antidote,  if  known,  should  be 
given.     (5)  If  irritating  gases  are  inhaled,  breathe  plenty  of 
fresh  air;   if  the  gases  get  into  the  eyes,  wash  the  eyes  freely 
with  water  and  then  drop  in  weak  boric  acid  or  borax  solution 
with  a  medicine  dropper.     (6)  Faintness  may  be  overcome 
by  holding  a  handkerchief  moistened  with  ammonia  or  cam- 


io  CHEMISTRY 

phor  near  the  nose.  (7)  Fires  may  be  extinguished  by  sand 
or  by  carbon  tetrachloride.  If  the  clothing  catches  fire,  a 
damp  towel  or  asbestos  blanket  should  be  used.  (8)  An 
emergency  box  or  cabinet  provided  with  the  following  articles 
should  be  kept  in  a  convenient  place:  Absorbent  cotton, 
bandages,  court  plaster,  pins,  thread,  scissors,  collodion, 
carron  oil,  sodium  bicarbonate,  vaseline,  smelling  salts,  cam- 
phor solution,  mustard,  boric  acid  solution,  medicine  dropper, 
and  a  handbook  of  first  aid  to  the  injured.  There  should 
also  be  available  a  fire  extinguisher,  a  box  of  sand  (including 
a  scoop),  and  a  blanket. 

11.  Laboratory  Notebooks.  —  A  neat  and  accurate .  record 
of  all  experiments  performed  by  the  pupil  should  be  made  in 
a  notebook  provided  for  this  purpose.  This  record  and  the 
form  in  which  it  may  be  kept  will  vary  with%  conditions.  It 
should  contain  at  least  the  following:  —  (i)  The  number  and 
title  of  each  experiment  and  the  date  of  performing.  (2)  A 
brief  account  of  each  experiment  in  such  a  form  that  the  ex- 
periment can  be  repeated  without  error  or  the  essential  parts 
subsequently  used.  (3)  Answers  to  all  questions  —  not 
merely  yes  or  no,  but  answers  in  which  the  question  itself  is 
involved.  (4)  A  simple  sketch  of  the  apparatus.  (5)  All 
numerical  data  involved  in  weighings  and  calculations.  (6) 
An  index. 


EXPERIMENTS 

PROPERTIES  —  CHANGES 

Experiment  1  —  Properties  of  Copper  Sulphate 

MATERIALS.  —  Copper  sulphate,  test  tubes  and  rack,  Bunsen  burner, 
test  tube  holder,  iron  nail,  ammonium  hydroxide,  barium  chloride 
solution. 

(a)  Examine  some  copper  sulphate  and  observe  its  proper- 
ties. What  is  its  physical  state,  i.e.  is  it  a  solid,  a  liquid,  or 
a  gas?  What  is  its  color?  Drop  a  small  piece  into  a  test 
tube  half  full  of  water;  is  it  heavier  or  lighter  than  water? 
Is  it  soluble  in  water?  Conclusive  evidence  regarding  its 
solubility  may  be  obtained  by  heating  the  test  tube.  If  you 
are  unfamiliar  with  the  method  of  heating  usually  employed 
in  a  chemical  laboratory,  proceed  as  follows:  Connect  the 
Bunsen  burner  with  one  end  of  the  rubber  tube  and  slip  the 
other  end  tightly  over  the  gas  outlet,  turn  on  the  gas  and  light 
it;  rotate  the  ring  at  the  base  of  .the  burner  until  the  flame  is 
colorless  or  faint  blue,  and  finally  adjust  the  gas  pressure  until 
the  flame  is  about  10  centimeters,  or  four 
inches,  high.  Attach  the  test  tube 
holder  to  the  test  tube  just  below  the 
lip  (Fig.  107),  put  the  lower  part  of  the 
test  tube  in  the  flame  and  move  the  test 
tube  slowly  up  and  down,  taking  care 
to  incline  it  slightly  and  to  heat  only 
the  part  that  contains  the  liquid;  if 
the  liquid  boils  too  vigorously,  the 
test  tube  should  be  removed  from  the 

flame   or  held   above,  it.     Continue  to    Fi^-  I07' TT^8t  Tube 

.,      ,  .  ,  and  Holder, 

heat  gently  until    there    is    conclusive 

evidence    of   the    solubility  of    the    copper    sulphate.     Does 
copper  sulphate  dissolve  readily  in  water?     Stand  the  test 


12  CHEMISTRY 

tube  in  the  test  tube  rack  to  let  the  liquid  cool,  or  cool  it 
by  holding  the  test  tube  in  a  stream  of  water. 

(b)  Determine  the  properties  of  copper  sulphate  exhibited 
when  different  substances  act  upon  it.     If  the  liquid  formed 
by  heating  copper  sulphate  and  water  is  not  uniformly  colored, 
pour  it  into  another  test  tube  and  then  back  again  into  the 
original  test  tube.     When  the  liquid  is  thoroughly  mixed, 
divide  it  into  three  equal  parts,  using  test  tubes  as  containers. 
Incline  one  test  tube,  carefully  slip  a  clean  iron  nail  into  it, 
and  let  the  test  tube  stand  in  the  rack  several  minutes.   Mean- 
while add  ammonium  hydroxide  solution  to  the  second  part 
until  the  test  tube  is  about  half  full,  and  mix  the  liquids  by 
shaking  or  by  stirring  with  a  glass  rod.      Add  a  little  barium 
chloride  solution  to  the  third  part,  shake  well,  and  let  this 
test  tube  also  stand  undisturbed  in  the  rack.     Examine  the 
three  test  tubes  in  succession.     Pour  the  liquid  out  of  the  first 
test   tube,  remove   and   examine   the   nail.     What   does   the 
deposit  resemble?     What  is  the  deposit?    If  you  are  in  doubt, 
compare  the  deposit  with  a  piece  of  copper.     The  deep  blue 
liquid  in  the  second  test  tube  contains  a  dissolved  substance 
which  is  formed  when  copper  sulphate  and  ammonium  hy- 
droxide act  upon  each  other.     Likewise  in  the  third  test  tube, 
the  white  substance  which  settles  to  the  bottom  is  formed 
from  the  sulphate  portion  of  the  copper  sulphate  when  copper 
sulphate  acts  upon  barium  chloride. 

(c)  Summarize  briefly  the  properties  of  copper  sulphate, 
dividing  them,  as  far  as  possible,  into  physical  and  chemical 
properties. 

Experiment  2  —  Physical  and  Chemical  Changes 

MATERIALS.  —  Small  piece  of  wood,   test  tubes  and  holder,  burner, 
fusible  metal,  glass  rod,  sulphur,  block  of  wood. 

A.  Wood.  Slip  a  small  piece  of  dry  wood  into  a  test  tube, 
attach  the  holder,  and  heat  cautiously;  hold  the  test  tube  so 
that  the  open  end  is  slightly  the  lower,  and  move  the  test 


EXPERIMENTS  13 

tube  slowly  back  and  forth  in  the  flame.  Heat  until  there  is 
definite  evidence  of  a  change  in  the  wood,  and  then  remove 
the  test  tube  from  the  flame.  Slip  out  the  solid,  and  when 
cool  examine  it.  Has  the  essential  change  in  the  wood  been 
physical  or  chemical? 

B.  Fusible  Metal.     Examine  a  small,  thin  piece  of  fusible 
metal  and  note  its  characteristic  properties.     Fill  the  test 
tube  half  full  of  water,  attach  the  holder,  and  heat  the  water 
to  boiling,  taking  care  not  to  heat  the  test  tube  above  the  sur- 
face of  the  water.    When  the  water  is  boiling,  remove  the  test 
tube  from  the  flame,  slip  the  metal  into  the  test  tube,  and 
observe  the  change  in  the  metal,  if  any.     Cool  the  water  by 
holding  the  lower  part  of  the  test  tube  in  a  stream  of  water. 
When  the  test  tube  is  cool  enough  to  handle  without  discom- 
fort, pour  off  the  water,  and  slip  out  the  solid.    Examine  the 
metal  carefully  and  compare  its  properties  with  those  origi- 
nally observed.     What    kind   of    a    change    did   the   metal 
undergo? 

C.  Glass  or  Rubber.     Rub  a  glass  rod  or  a  fountain  pen 
briskly  on  a  piece  of  cloth,  and  hold  it  near  very  small  bits 
of  dry  paper.    Describe  the  result.    After  a  moment  try  again. 
What  kind  of  a  change  did  the  glass  or  the  rubber  (of  the  pen 
holder)  undergo? 

D.  Sulphur.     Examine  a  piece  of  sulphur  and  note  its  prop- 
erties,  e.g.   color,    brittleness,   solid    condition.     Put  a  small 
piece  on  a  block  of  wood  and  light  the  sulphur  by  directing 
the  flame  upon  it.    Observe  the  color  and  size  of  the  flame  of 
the  burning  sulphur.    Observe  also  (very  cautiously)  the  odor 
of  the  gaseous  product  by  wafting  a  little  gently  toward  the 
nose.    Compare  the  properties  of  this  gaseous  substance  with 
those  of  the  sulphur.    Has  the  essential  change  in  the  sulphur 
been  physical  or  chemical?    Why? 

NOTE.  —  As  soon  as  the  properties  of  the  burning  sulphur  have  been 
observed,  extinguish  it  with  a  little  sand  or  by  pressing  it  with  a  piece 
of  stiff  paper. 


14  CHEMISTRY 

SUPPLEMENTARY  EXPERIMENTS 

Not  all  the  Supplementary  Experiments  need  be  done.  Those 
should  be  selected  that  are  needed  to  emphasize  certain  applications  or 
principles.  These  Experiments  may  also  be  assigned  to  pupils  who 
work  quickly  or  who  need  special  preparation  for  examinations. 


Experiment  3  —  Properties  of  Iodine 

MATERIALS.  —  Iodine,  alcohol,  carbon  disulphide. 

(a)  Examine  a  piece  of  iodine  and  observe  its  physical  state,  luster, 
color,  and  odor.     Touch  it  with  the  finger  and  observe  the  effect 
upon  the  skin.    Drop  a  piece  into  a  test  tube  half  full  of  water.    Is 
it  heavier  or  lighter  than  water?    Stand  the  test  tube  in  the  rack  and 
let  it  remain  undisturbed  until  needed  for  (c). 

(b)  Drop  a  piece  of  iodine  into  a  dry  test  tube,  grasp  the  test  tube 
near  the  top  with  the  test  tube  holder,  and  gently  heat  the  bottom 
of  the  test  tube  in  the  upper  part  of  the  flame  until  a  definite  change 
occurs  in  the  iodine.    Whit  is  the  effect  of  heat  upon  iodine? 

(c)  The  solubility  of  iodine  may  now  be  determined.    Shake  the 
test  tube  containing  the  water  and  iodine,  let  any  undissolved  iodine 
settle  and  then  pour  the  liquid  into  another  test  tube.    Examine  this 
liquid.     Is  there  evidence  of  dissolved  iodine?     If  the  evidence  is 
inconclusive,  add  a  few  drops  of  carbon  disulphide  and  shake  well. 
Carbon  disulphide  is  much  heavier  than-  water  and  sinks  to  the  bot- 
tom; at  the  same  time  it  absorbs  any  dissolved  iodine  and  becomes 
violet  in  color.    What  final  conclusion  can  be  drawn  regarding  the 
solubility  of  iodine  in  water?    Measure  1 5  cubic  centimeters  of  alcohol 
in  the  graduated  cylinder,  and  pour  it  into  the  other  test  tube  that 
contains  the  piece  of  undissolved  iodine.   Shake  well,  and  warm  slightly 
by  holding  the  test  tube  above  a  low  flame  for  a  minute  or  two; 
take  care  not  to  set  the  alcohol  on  fire.    Shake  well.    What  is  the  evi- 
dence of  the  solubility  of  iodine?    Confirm  the  conclusion  by  pouring 
a  little  of  this  liquid  into  a  test  tube  half  full  of  water,  adding  a  few 
drops  of  carbon  disulphide,  and  shaking  well.    What  final  conclusion 
can  now  be  drawn  regarding  the  solubility  of  iodine  in  alcohol? 

(d)  Summarize  briefly  the  properties  of  iodine. 


EXPERIMENTS  15 

Experiment  4  —  Physical  and  Chemical  Changes 

MATERIALS.  —  Copper  wire,  electric  bell  apparatus. 

(a)  Examine  a  piece  of  clean  copper  wire  and  notice  especially 
its  color  and  flexibility.  Grasp  one  end  of  the  wire  with  the  forceps, 
and  hold  the  other  end  in  the  hottest  part  of  the  flame  until  the  cop- 
per melts  and  undergoes  a  definite  change.  Then  remove  it  from  the 
flame  and  examine  the  black  product.  Compare  its  properties  with 
those  of  copper.  Is  it  apparently  a  different  substance  from  the 
copper?  Why?  What  kind  of  a  change  did  the  melted  copper  undergo? 

(b~)  Introduce  a  piece  of  copper  wire  into  the  circuit  of  an  electric 
bell  apparatus.  Does  copper  conduct  electricity?  Remove  the  wire 
and  examine  it.  What  kind  of  a  change  did  it  undergo? 

(c)  Roll  a  piece  of  copper  wire  into  a  ball,  drop  it  into  a  test  tube 
half  full  of  dilute  nitric  acid,  and  warm  gently.  What  is  the  evidence 
that  the  copper  is  undergoing  a  change?  Verify  the  observation  by 
utilizing  a  preceding  experiment.  What  kind  of  a  change  did  the 
copper  undergo  when  treated  with  nitric  acid? 

Experiment  5  —  Physical  and  Chemical  Changes 

MATERIALS.  —  Magnesium,  forceps. 

Examine  a  piece  of  magnesium  and  note  its  properties,  especially 
the  luster,  color,  and  flexibility.  Grasp  one  end  firmly  with  the  forceps 
and  hold  the  other  end  in  the  flame  for  an  instant  and  then  remove 
it.  Observe  the  result.  Examine  the  whitish  substance  that  is  formed 
and  compare  its  properties  with  those  of  magnesium.  Has  the 
essential  change  in  the  magnesium  been  physical  or  chemical?  Why? 


OXYGEN 
Experiment  6  —  Preparation  and  Properties  of  Oxygen 

MATERIALS.  —  15  grams  of  potassium  chlorate,  15  grams  of  manganese  diox- 
ide, 5  bottles  (about  250  cubic  centimeters  each),  filter  paper,  joss 
stick  or  splint  of  wood,  sulphur,  deflagrating  spoon,  piece  of  charcoal 
fastened  to  one  end  of  a  copper  wire  (30  centimeters  long)  and  a  wad 
of  iron  thread  (often  called  "  steel  wool " )  to  the  other  end.  The 
apparatus  is  shown  in  Fig.  108.  A  is  a  large  test  tube  provided  with 
a  one-hole  rubber  stopper,  to  which  is  fitted  a  short  glass  tube  B;  the 
delivery  tube  D  is  attached  to  the  short  glass  tube  by  the  rubber 
tube  C. 

I.  Preparation.  Weigh  the  potassium  chlorate  on  a  piece 
of  paper  creased  lengthwise,  and  slip  it  into  the  test  tube;  do 
the  same  with  the  manganese  dioxide.  Shake  the  test  tube 
until  the  chemicals  are  thoroughly  mixed;  then  hold  the  test 
tube  in  a  horizontal  position  and  roll  or  shake  it  until  the  mix- 
ture is  spread  along  about  one  half  of  the  tube.  Insert  the 
stopper  with  its  tubes,  and  clamp  the  test  tube  to  the  iron 
stand,  as  shown  in  Fig.  108,  taking  care  not  to  crush  the  tube. 

Fill  the  pneumatic  trough  with  water,  until  the  support  is 
just  covered.  Fill  the  bottles  full  of  water,  cover  each  with  a 
piece  of  filter  paper,  invert  one  of  them  in  the  trough,  remove 
the  filter  paper,  and  stand  the  inverted  bottle  upon,  or  near,  the 
support.  The  end  of  the  delivery  tube  D  should  rest  on  the 
bottom  of  the  trough,  just  under  the  hole  in  the  support. 

Before  proceeding,  ask  the  Teacher  to  inspect  the  apparatus. 
Heat  the  whole  test  tube  gently  with  a  flame  about  10  cen- 
timeters (or  4  inches)  high.  When  the  gas  bubbles  regularly 
through  the  water  slip  the  inverted  bottle  over  the  hole  in  the 
support.  The  gas  will  rise  in  the  bottle  and  force  out  the 
water.  Move  the  flame  slowly  along  the  test  tube,  taking  care 
not  to  heat  the  tube  too  long  in  one  place  nor  too  near  the 
.rubber  stopper.  If  the  gas  is  evolved  too  rapidly,  lessen  the 


OXYGEN  17 

heat;  if  too  slowly,  increase  it;  if  not  at  all,  examine  the  stopper 
and  the  rubber  connecting  tube  for  leaks,  and  adjust  accord- 
ingly. When  the  first  bottle  of  gas  is  full,  remove  it,  cover 
it  with  a  piece  of  wet  filter  paper,  and  stand  it  upon  the  desk; 
invert  another  bottle,  remove  the  filter  paper,  and  slip  the 
bottle  over  the  hole.  When  five  bottles  of  gas  have  been 
collected,  immediately  remove  the  end  of  the  delivery  tube 
from  the  water,  lest  the  cold  water  be  drawn  up  into  the  hot 
test  tube  as  the  gas  contracts.  Perform  II  at  once. 

II.  Properties.  Proceed  as  follows  with  the  oxygen  prepared 
in  I.  (a)  Dip  a  glowing  joss  stick  into  one  bottle,  and  observe 
the  change.  Remove  the  joss  stick,  make  it  glow  again,  and  re- 
peat as  many  times  as  possible.  How  does  the  glowing  joss 
stick  change?  Does  the  oxygen  burn?  What  property  of 
oxygen  does  this  experiment  show? 

(b)  Put  a  small  piece  of  sulphur  in  the  deflagrating  spoon, 
hold  the  spoon  in  the  Bunsen  flame  until  the  blue  flame  of  the 


Fig.  108.  —  Apparatus  for  Preparing  Oxygen. 

burning  sulphur  can  be  seen,  then  lower  the  spoon  into  a  bottle 
of  oxygen.  Notice  any  change  in  the  flame.  Waft  a  little  of 
the  gaseous  product  toward  the  nose.  Of  what  does  the  odor 
remind  you?  As  soon  as  the  results  are  conclusive,  remove  the 
spoon  and  plunge  it  into  the  water  in  the  trough  to  extinguish 
the  burning  sulphur. 


1 8  CHEMISTRY 

(c)  Hold  the  charcoal  in  the  flame  long  enough  to  produce 
a  faint  glow,  then  lower  it  into  a  bottle  of  oxygen.  Observe 
the  result. 

(d)   Twist  one   end  of  the  copper  wire  (used 
in  (c))  firmly  around  the  wad  of  iron  thread 
(Fig.  109),  heat  the    ends  of  a  few  strands  of 
the  thread  an  instant  in  the  flame,  and  quickly 
lower   it   into   a   bottle    of    oxygen.    The   iron 
thread   should  change  conspicuously.     If  it  does 
Fig.  109.— Iron  not,  heat  it  a  second  time  in  the  flame,  and  lower 
SchTcfto  End  ifc  again  into  tne  bottle  of  oxygen.     Observe  the 
of    Copper  result. 

Wire-  (e)   With  the  remaining  bottle,  repeat  any  of 

the  above  experiments. 
NOTE.  — •  Clean  the  test  tube  used  in  6  I  with  a  little  warm  water. 

Required  Exercises. —  i.  Write  a  brief  account  of  Exp.  6  I  in  your  note 
book. 

2.  Write  a  brief  account  of  Exp.  6  II,  answering  all  questions. 

3.  Sketch  the  apparatus  used  to  prepare  oxygen. 

Experiment  7  —  Oxidation  of  Copper 

MATERIALS.  — •  Copper  borings,  evaporating  dish,  gauze-covered  ring,  iron 
stand,  test  tube  and  cork. 

Put  about  4  gm.  of  copper  borings  in 
an  evaporating  dish  and  stand  the  dish  on 
a  gauze-covered  ring,  which  is  attached  to 
an  iron  stand  (Fig.  no).  Heat  the  dish 
carefully  but  strongly  about  ten  minutes; 
then  direct  the  free  flame  of  the  burner  upon 
the  contents  of  the  dish  for  about  five 
minutes,  stirring  occasionally  with  a  glass 
rod.  Describe  any  marked  change  in  the 
copper.  When  the  dish  is  cool,  pour  the 

contents  into  a  test  tube,  cork  the  test  tube  Fi     IIQ E  v  a  p  o  - 

tightly,  and  save  for  use  in  Exp.  11.  rating  Dish  on  a 

Describe  the  experiment  briefly.     What     Gauze-covered  Ring, 
chemical  compound  was  formed  ?     What  elements  combined  ? 


OXYGEN 


What  general  name  is  given  to  this  kind  of  chemical  change  ? 
What  spec'al  name  ? 

NOTE.  —  The  dish  can  be  cleaned  by  warming  dilute  nitric  acid  in  it. 

SUPPLEMENTARY  EXPERIMENTS 

(See  note  on  page  14.) 

Experiment  8  —  Preparation  of  Oxygen  from  Various 
Substances 

(Each  pupil  need  not  perform  all  of  this  experiment.} 

MATERIALS.  —  Mercuric  oxide,  lead  dioxide,  barium  dioxide,  sodium  perox- 
ide, hydrogen  peroxide,  potassium  permanganate,  joss  stick. 

A.  Mercuric  Oxide.     Put  a  little  mercuric  oxide  on  the  end  of  a 
narrow  piece  of  paper  creased  lengthwise,  and  slip  the  powder  into  a 
test  tube.     The  powder  should  nearly  fill  the  round  end  of  the  test 
tube.     Hold  the  test  tube  in  a  horizontal  posi- 
tion, shake  it  to  spread  the  powder  into  a  thin 

layer,  and  then  clamp  the  test  tube  in  the 
position  shown  in  Fig.  in.  Heat  the  test 
tube  strongly  with  the  upper  part  of  the 
Bunsen  flame.  Do  not  heat  one  place,  but 
move  the  burner  back  and  forth.  As  soon  as 
a  definite  change  is  noticed  inside  the  tube, 
insert  a  glowing  joss  stick.  Observe  and  de- 
scribe the  change.  If  there  is  no  change,  heat 
strongly,  and  test  again.  What  gas  is  liber- 
ated? Examine  the  deposit  inside  the  tube. 
What  is  it?  If  you  are  in  doubt,  scrape  out 
a  little,  and  examine  again.  State  the  result 
of  the  final  observation. 

B.  Lead  Dioxide.     Put  a  little  lead  dioxide 


Fig.  in.— Test  Tube 
Clamped  in  Position 
for  Heating  Certain 
Substances. 


into  another  test  tube  and  proceed  with  the  heating  as  in  A.     Test 
with  the  joss  stick.     Observe  and  state  the  result. 

C.  Barium  Dioxide.  Proceed  as  in  B  using  barium  dioxide. 

D.  Sodium  Peroxide  and  Water.   Fill  a  test  tube  two-thirds  full  of 
water  and  stand  it  in  the  test  tube  rack.     Obtain  from  the  Teacher 
a  little  sodium  peroxide  on  a  creased  paper,  cautiously  slip  the  sodium 
peroxide  into  the  water,  and  then  thrust  a  glowing  joss  stick  into  the 
upper  part  of  the  test  tube.     Observe  and  state  the  result. 


20  CHEMISTRY 

E.  Hydrogen  Peroxide  and  Potassium  Permanganate.  Fill  a  test 
tube  one-third  full  of  hydrogen  peroxide,  add  half  the  volume  of  dilute 
sulphuric  acid,  and  then  several  drops  of  potassium  permanganate 
solution.  Test  as  in  D.  State  the  result. 

Experiment  9  —  Effect  of  Heating  a  Known  Weight  of  a 
Metal  in  Air 

MATERIALS.  —  Porcelain  crucible  and  cover,  zinc  dust,  triangle,  scales. 

Clean  and  dry  a  porcelain  crucible  and  cover,  and  weigh  both  on 
the  scales  (or  on  the  balance,  if  desired).  Record  the  weight  as  shown 
below.  Crease  a  slip  of  paper  lengthwise,  pour  zinc  dust  into  the 
crease  and  slide  the  zinc  dust  into  the  crucible  until  about  3  gin. 
have  been  added,  and  then  weigh  accurately  (including  the  cover). 
Record  as  below.  Place  the  covered  crucible  on  the  triangle,  which 
may  be  supported  by  a  ring  attached  to  an  iron  stand.  Heat  gently 
with  a  low  flame  to  avoid  breaking  the  crucible.  Gradually  increase 
the  heat  until  the  flame  is  just  above  the  bottom  of  the  crucible. 
Heat  for  about  twenty  minutes.  Lift  the  cover  occasionally  by  grasp- 
ing the  ring  with  the  forceps.  If  the  zinc  glows  and  a  smoke  escapes, 
cover  the  crucible  at  once  to  prevent  loss;  it  is  necessary  to  admit 
air  and  to  heat  the  zinc  very  hot,  but  little  or  nothing  should  be  al- 
lowed to  escape  from  the  crucible.  Cool  the  crucible  gradually  by 
moving  the  flame  slowly  beneath  it.  As  soon  as  the  crucible  is  cool, 
weigh,  and  record  as  below.  To  what  is  the  change  in  weight  due? 

RECORD 

Weight  of  crucible,  cover,  and  zinc 

and  cover 

"       "  zinc 

"       "  crucible  and  contents  before  heating  . 

"       "         "         "  "       after          "      .. 

Change  in  weight    

NOTE. —  The  crucible  can  be  cleaned  with  dilute  hydrochloric  acid. 


HYDROGEN 
Experiment  10  —  Preparation  and  Properties  of  Hydrogen 

MATERIALS. —  Granulated  zinc,  dilute  sulphuric  acid,  four  bottles,  filter 
paper,  taper,  matches,  and  the  apparatus  shown  in  Fig.  112.  A  is 
a  bottle  provided  with  a  two-hole  stopper,  through  which  passes 
the  dropping  tube  B  and  the  right-angle  bend  C;  the  (15  centi- 
meters or  6  inches)  tube  D  is  attached  to  the  bent  tube  by  the  rubber 
tube  E.  The  dropping  tube  is  made  as  follows:  Cut  off  the  top  of  a 
thistle  tube  about  2.5  centimeters  (i  inch)  below  the  juncture  of 
the  stem  and  cup,  and  heat  the  sharp  ends  a  minute  or  two  in 
the  flame;  when  cool,  slip  a  thick-walled  rubber  tube  (5  centimeters 
or  2  inches  long)  over  one  end  of  the  stem,  attach  a  pinch-clamp 
to  the  rubber  tube,  and  connect  the  tube  with  the  cup,  taking  care 
to  have  the  ends  of  the  glass  tubes  as  near  together  as  possible; 
if  properly  constructed,  the  cup  will  remain  upright  when  full  of 
liquid. 

I.  Preparation.  Weigh  about  40  gm.  of  granulated  zinc 
and  slip  it  into  the  bottle.  Insert  the  stopper  with  its  tubes. 
Fill  the  pneumatic 
trough  with  water 
as  usual,  and  adjust 
the  apparatus  so 
that  the  end  of  the 
delivery  tube  rests 
on  the  bottom  of 
the  trough  under 
the  hole  in  the  sup- 
port. Fill  the  bottles 
with  water,  and 
cover  with  filter 
paper;  invert  one 

in  the   trough,    re- 
Fig.  112.  —  Apparatus  for  Preparing  Hydrogen, 
move     the     paper, 

and  stand  the  inverted  bottle  upon  the  support. 

Fill  the  cup  with  dilute  sulphuric  acid,  and  let  the  acid  run 


22  CHEMISTRY 

into  the  bottle  by  pinching  the  clamp;  if  the  acid  does  not 
flow  freely  down  the  tube  into  the  bottle,  loosen  the  stopper 
for  an  instant,  and  as  soon  as  the  acid  enters  the  bottle  push 
the  stopper  into  place.  The  interaction  of  the  zinc  and  sul- 
phuric acid  produces  hydrogen,  and  the  gas  will  bubble 
through  the  water  in  the  trough  up  into  the  bottle.  Collect 
and  remove  the  bottles  of  gas  as  in  the  Preparation  of  Oxygen, 
taking  care  to  cover  each  bottle  tightly  with  a  piece  of  wet 
filter  paper.  If  the  evolution  of  gas  slackens  or  ceases,  add  a 
little  more  acid  through  the  dropping  tube.  Collect  four 
bottles  of  hydrogen,  and  perform  II  at  once. 

II.  Properties.  Proceed  as  follows  with  the  hydrogen  gas 
prepared  in  I.  (a)  Uncover  a  bottle  for  an  instant  to  let  a 
little  air  in,  and  then  drop  a  lighted  match  into  the  bottle. 
Observe  the  result. 

(b)  Remove  the  paper  from  another  bottle  of  hydrogen, 
and  allow  it  to  remain  uncovered  for  three  minutes  —  by  the 
clock.     Then  show  the  presence  or  absence  of  hydrogen  by 
dropping  a   lighted    match   into    the    bottle.      Observe   the 
result.      What  property  of  hydrogen  is    shown  by  this  ex- 
periment? 

(c)  Verify  your  answer  to  the  last  question,  thus:    Hold  a 

bottle  of  air  over  a'  covered  bottle  of  hydro- 
gen, remove  the  paper,  and  bring  the  mouths 
of  the  bottles  together,  as  shown  in  Fig.  113. 
Hold  them  there  for  a  minute  or  two,  then 
stand  the  bottles  on  the  desk  and  cover  them 
with  wet  filter  paper.  Drop  a  lighted  match 
into  each  bottle.  Observe  the  result.  How 
does  (c)  verify  (b)? 

(d)  Invert  a  covered  bottle  of  hydrogen,  re- 
move the  paper,  and  quickly  thrust   a  lighted 
Fig.  113—         taper  up  into  the  bottle.     Withdraw  the  taper, 
Transferring       then  insert  and  withdraw  it  several  times,  and 
observe   carefully    (i)    whether    the    hydrogen 
burns,  and,  if  so,  where?  and    (2)  if  the  taper  burns  inside 


HYDROGEN 


and  outside   the   bottle, 
scribe  and  explain. 


Feel  of  the  neck  of  the  bottle;  de- 


NOTE.  —  As  soon  as  this  experiment  is  completed,  pour  off  the  acid 
from  any  unused  zinc.  If  Exp.  11  is  not  to  be  performed  soon,  wash 
the  zinc  several  times,  and  save  it  for  other  experiments. 

Required  Exercises.  —  i.  Write  a  brief  account  of  Exp.  10  I. 

2.  Write  a  brief  account  of  Exp.  10  II,  answering  all  questions. 

3.  The  apparatus  used  to  prepare   hydrogen  is  called  a  generator; 
sketch  it  (from  memory,  if  possible). 

Experiment  11  —  Reduction  of  Copper  Oxide  by  Hydrogen 

MATERIALS.  —  Copper  oxide,  apparatus  shown  in  Fig.  114. 

Put  the  copper  oxide  that  was  prepared  in  Exp.  7  in  an 
evaporating  dish,  stand  the  dish  on  a  gauze-covered  ring 
attached  to  an  iron  stand  (Fig.  no),  and  heat  gently.  Mean- 


ir 


F 


\ 


Fig.  114.  —  Apparatus  for  the  Reduction  of  Copper  Oxide 
by  Hydrogen. 

while,  arrange  the  apparatus.  The  parts  lettered  A,  B,  C, 
D,  E,  constitute  the  hydrogen  generator  and  are  the  same  as 
those  used  in  Exp.  10  I.  F  is  a  large  test  tube  fitted  with  a 
two-hole  stopper;  the  delivery  tube  E  passes  through  one 


24  CHEMISTRY 

hole  and  extends  nearly  to  the  bottom,  and  the  right-angle 
tube  G  passes  just  through  the  other  hole;  the  tube  G  is 
lengthened  by  the  rubber  tube  H. 

Slip  the  copper  oxide  into  the  test  tube  F,  hold  the  test  tube 
in  a  horizontal  position  and  tap  it  gently  to  spread  the  solid 
into  a  long,  thin  layer.  Connect  the  test  tube  with  the  rest 
of  the  apparatus,  and  clamp  it  into  the  proper  position,  taking 
care  not  to  crush  the  tube. 

Ask  the  Teacher  to  inspect  the  apparatus,  and  do  not  pro- 
ceed until  permission  is  given.  After  obtaining  permission, 
fill  the  cup  of  the  generator  nearly  full  with  dilute  sulphuric 
acid,  pinch  the  clamp,  and  let  about  half  the  acid  run  into  the 
generator  bottle.  Allow  the  gas  to  flow  steadily  for  at  least 
two  minutes  before  lighting  the  Bunsen  burner;  then  intro- 
duce a  little  more  acid.  Now,  heat  gently  the  lower  part  of 
the  test  tube  where  the  copper  compound  is  located.  Do  not 
let  the  flame  come  near  the  rubber  tube  H.  The  gas  must  flow 
slowly  through  the  apparatus  during  the  heating;  if  it  does 
not  (as  you  can  tell  by  the  bubbles  in  the  bottle  or  by  smelling 
the  gas  at  the  end  of  the  rubber  exit  tube),  introduce  a  little 
more  acid.  If  the  test  tube  F  should  break,  pinch  the  rubber 
tube  D  an  instant  to  cut  off  the  flow  of  hydrogen,  and  then 
extinguish  the  Bunsen  burner  flame.  When  a  marked  and 
permanent  change  is  observed  inside  the  test  tube  F,  stop 
heating,  and  extinguish  the  Bunsen  burner  flame  at  once. 
Observe  the  entire  contents  of  the  test  tube;  what,  in  all 
probability,  are  both  products? 

Required  Exercises. —  i.  Describe  briefly  the  whole  experiment, 
and  sketch  the  apparatus. 

2.  What  chemical  compound  was  formed  in  F? 

3.  How  was  the  copper  compound  changed?     What  special  name 
is  given  to  this  kind  of  a  change? 

4.  Recall  Exp.  7;   in  Exps.  7  and  11,  what  was   oxidized,  what  was 
reduced,   and   what  substances  accomplished   the   oxidation  and   the 
reduction?     Summarize  in  a  few  words  how  Exps.  7  and  11  illustrate 
oxidation  and  reduction. 


HYDROGEN 


SUPPLEMENTARY  EXPERIMENTS 
(See  note  on  page  14.) 

Experiment  12  —  Preparation  of  Hydrogen  from  Various 
Compounds 

(//  desired,  D  and  E  may  be  performed  later  as  Exp.  29  A.) 

MATERIALS.  —  Magnesium,  aluminium,  zinc,    iron,  hydrochloric  acid 
sulphuric  acid,  sodium  hydroxide,  sodium,  calcium. 

A.  Metals  and  Hydrochloric  Acid.    Fill  a  test  tube  half  full  of  dilute 
hydrochloric  acid,  stand  the  test  tube  in  the  rack,  and  drop  in  a  small 
piece  of  magnesium.     In  a  minute  or  two  test  the  escaping  gas  by 
holding  a  lighted  match  (or  a  low  flame)  at  the  mouth  of  the  test  tube. 
What  gas  is  it?    What  was  its  source?     (If  the  test  is  not  decisive, 
add  more  magnesium,  or  wait  until  more  gas  accumulates  in  the  test 
tube.)     Proceed  in  the  same  way  with  aluminium,  zinc,  and  iron 
(in  the  form  of  filings);    use  separate  test  tubes,  and  heat,  if  the 
action  is  slow.     Observe  the  result  in  each  case,  and  apply  the  ques- 
tions asked  about  magnesium. 

B.  Metals  and  Sulphuric  Acid.    Proceed  as  in  A  and  observe 
the  result  in  each  case.    Answer  the  questions  asked  in  A. 

C.  Aluminium    and    Sodium    Hydroxide. 
Roll  two  or  three  pieces  of  aluminium  into 
a  ball,    drop  it   into   a   test  tube,   slip  in 
a  piece  of  sodium  hydroxide  about  2.5  cm. 
(or    i    in.)    long,  and   add    a  little   water. 
Warm    slightly,   if   no    action    results,  and 
test    as  above.     Observe  the   result, 
swer  the  questions  asked  in  A. 

D.  Sodium  and  Water.    Precaution. 
dium   should    be    handled    cautiously    and 
used  strictly  according  to  directions.     Small 
fragments  obtained  for  experiments  should 
be  protected  by  a  mortar  or  dish.      If  so- 
dium is  left  from   an  experiment,  it   must 
not    be   thrown    into    the   refuse    jar,    but 
returned  to  the  Teacher.  —  Fill  a  porcelain 
evaporating  dish  two-thirds  full  of  water. 


An- 


Fig.  115. —  Apparatus 
for  collecting  the  Gas 
Liberated  by  the  In- 
teraction of  Water 
and  Metals. 


FiU  a  small  test  tube 


26  CHEMISTRY 

full  of  water,  cover  and  invert  it,  and  clamp  it  as  shown  in  Fig.  115. 
Wrap  a  small  piece  of  clean  sodium  loosely  in  a  piece  of  tea  lead 
about  5  centimeters  (2  inches)  square,  make  two  or  three  small  holes 
in  the  tea  lead,  and  then  thrust  it  under  the  test  tube.  A  gas  will 
rise  into  the  test  tube.  Proceed  similarly  with  additional  pieces  of 
sodium  and  dry  tea  lead  until  the  test  tube  is  full  of  gas;  then 
unclamp  it,  remove,  and  invert.  Hold  a  lighted  match,  for  an  in- 
stant, at  the  mouth  of  the  tube.  Observe  the  result,  especially  at 
the  mouth  of  the  tube.  What  is  the  gas?  What  was  it  source? 

E.  Calcium  and  Water.  Arrange  the  apparatus  as  in  D.  Drop 
two  or  three  pieces  of  calcium  into  the  water  in  the  dish,  and  push 
them  under  the  test  tube.  As  soon  as  the  tube  is  full,  or  nearly 
full,  remove  the  tube,  and  test  the  gas  as  in  D.  State  the  result. 
Answer  the  questions  asked  in  D. 


PROPERTIES   OF    GASES 

SUPPLEMENTARY  EXPERIMENT 
Experiment  13  —  Weight  of  a  Liter  of  Oxygen 

(If  desired,  this  Experiment  may  be  postponed  until  the  pupil  has 
acquired  more  experience  in  the  laboratory.  It  is  suggested  that  it  be 
performed  while  Chapter  VII  or  VIII  is  being  studied.) 

MATERIALS.  —  Potassium  chlorate,  manganese  dioxide,  calcium  chloride, 
glass  wool  or  shredded  asbestos,  and  the  apparatus  shown  in  Fig. 
1 1 6.  A  is  a  large  test  tube  attached  to  the  bent  tube  F  by  a  rub- 
ber stopper.  B  is  a  large  bottle  to  be  filled  with  water;  it  is  pro- 
vided with  a  two-hole  rubber  stopper,  through  which  pass  F  and  C, 
the  latter  being  connected  with  a  rubber  tube  C'  to  which  is 
attached  the  short  glass  tube  G.  A  Hofmann  screw  is  attached  at 
the  point  E.  Another  large  bottle  D  serves  to  catch  the  water 
forced  over  from  B  through  CC  by  the  oxygen  generated  in  A.  A 
hook  (S)  of  aluminium  wire  permits  A  to  be  hung  from  the  balance 
beam  in  weighing. 

Fill  the  space  i  in  A  with  a  mixture  of  manganese  dioxide  and 
potassium  chlorate  (about  equal  parts).  Each  substance  must  be 
powdered  and  free  from  organic  matter  (e.g.  paper,  cork,  straw). 
The  mixture  should  be  dried  by  heating  it  in  an  oven  to  about  110° 
C.,  on  a  radiator,  or  on  some  convenient  heated  object.  Push  glass 
wool,  or  shredded  asbestos  (previously  ignited  to  a  red  heat),  into 
the  space  2  in  A.  Put  small  lumps  of  calcium  chloride  into  3  and 
glass  wool  into  4.  Push  the  stopper  well  into  the  test  tube.  Wipe  A 
carefully  with  soft  paper,  and  then  weigh  AF  on  the  balance.  Weigh 
the  empty,  dry,  clean  bottle  D  to  a  decigram  on  the  scales. 

Fill  B  with  water  nearly  to  the  neck.  Fill  CC'  with  water  and 
tighten  the  Hofmann  screw  to  prevent  the  water  from  running  out. 
Insert  F  into  the  stopper  of  B,  Push  the  stopper  into  the  bottle, 
slowly  at  first,  then  hard;  if  water  rises  in  F,  loosen  the  screw  at  E 
slightly,  remove  A ,  and  blow  gently  into  F  to  force  the  water  back 
into  B.  When  properly  adjusted,  the  water  should  be  in  B  and  CC', 
but  not  in  F.  Replace  A ,  taking  care  not  to  crush  the  thin  glass  by 


28 


CHEMISTRY 


pushing  it  too  hard  upon  its  stopper;   open  the  screw  at  E.    If  the 

apparatus  is  tight,  little  or  no  water  will  flow  out.     It  should  be 

adjusted  until  air  tight.    Leave  the  screw  open. 

Heat  A  gently  with  a  low  flame,  keeping  the  flame  back  of  the 

space  2.    The  liberated  oxygen  will  force  the  water  from  B  into  D. 

Heat  A  just  hot  enough  to  cause  a  gentle  flow  of  water  into  D.    When 

D  is  about  three-fourths  full, 
decrease  the  heat  gradually. 
While  A  is  cooling  sufficiently 
to  weigh,  stand  a  thermometer 
in  D;  also  read  the  barometer. 
When  A  is  cold,  raise  B  until 
the  water  is  at  the  same  level 
in  B  and  Z>,  pinch  C"  tight  and 
remove  it  from  D.  Read  and 
remove  the  thermometer.  Dry 

Fig.  1 16.-  Apparatus  for  Finding  the   D  on  the  outside>  if  necessary, 
Weight  of  a  Liter  of  Oxygen.          and   then  weiSh   *>    using    the 

same  large  weights  as   before; 

the  gain  in  weight  (in  grams)  of  D  gives  the  volume  (since  i  gm.  of 
water  =  i  cc.)  of  oxygen  liberated.  Weigh  AF;  its  loss  in  weight  is 
the  weight  of  the  oxygen  that  passed  into  B. 

Reduce  the  observed  volume  to  the  volume  it  would  occupy,  if 
it  were  at  o  C.,  760  mm.,  and  in  the  dry  state.  This  is  done  by  the 
formula  — 

V'  (P'  -  a) 


V  = 


760(1+ (.00366X1)) 


Substitute  the  proper  values  in  this  formula,  and  solve  for  V  —  the 
corrected  volume  of  oxygen  liberated.      (See  Part  I,  33,  34,  40.) 

Since  i  liter  contains  1000  cubic  centimeters,  then  V/iooo  is  the 
actual  volume  of  liberated  oxygen  expressed  in  liters.  The  weight 
of  liberated  oxygen  is  found  by  subtracting  the  weight  of  AF  after 
heating  from  its  weight  before  heating.  And  finally  the  weight  of 
i  liter  of  oxygen  in  grams  is  found  by  dividing  the  weight  of  liberated 
oxygen  by  its  volume. 


PROPERTIES   OF   WATER 
Experiment  14  —  Water  in  Food 

MATERIALS.  —  Substances  enumerated  below. 

Heat  gently  in  a  dry  test  tube  a  small  piece  of  meat.  Hold 
the  open  end  of  the  test  tube  lower  than  the  closed  end,  and 
take  care  not  .to  burn  the  substance.  What  substance  is 
liberated?  Repeat,  using  a  dry  test  tube  in  each  case  and  a 
small  piece  of  one  or  more  of  the  following:  Potato,  apple, 
cranberry,  celery,  bread,  cracker.  Observe  and  state  the 
result  in  each  case. 

Experiment  15  —  Some  Physical  Properties  of  Water 

MATERIALS.  —  Copper  wire,  ice,  thermometer,  salt. 

A.  Conduction  of  Heat.     Wind  enough  copper  wire  around  a 
small  lump  of  ice  to  make  it  sink  in  water,  slip  it  into  a  large 
test  tube  nearly  full  of  water,  and  quickly  heat  the  water  to 
boiling  near  the  surface.     Observe  the  result  as  soon  as  the 
water  boils.     What  does    this  experiment   show  about   the 
conducting  power  of  water? 

B.  Expansion  and  Contraction.     Fill  a  large  test  tube  full  of 
water,  and  insert   a  one-hole  rubber  stopper  fitted  with   a 
short  glass  tube.      Attach  the  test  tube  holder  and  heat  the 
water   slowly.     Observe  any  change    in   the  volume  as  the 
water  increases  in  temperature.     Now  cool  the  water  by  hold- 
ing the  test  tube  in  a  stream  of  running  water  and  observe  any 
change  in  the  volume.     What    does   this   experiment   show 
about  the  expansion  and  contraction  of  water  when  its  tem- 
perature changes? 

C.  Boiling  Point.     Fill  a  large  test  tube  half  full  of  water, 
clamp  it  in  an  upright  position  to  an  iron  stand,  and  heat  the 
water  to  boiling.    Hold  the  bulb  of  a  thermometer  in  the  escap- 
ing steam  and  note  the  highest  temperature  recorded.    Slowly 


30  CHEMISTRY 

lower  the  thermometer  until  the  bulb  touches  the  boiling  water, 
note  the  highest  temperature,  and  then  remove  the  ther- 
mometer. Compare  the  two  maximum  readings.  Average 
them  and  record  the  result. 

D.  Freezing  Point,  (a)  Put  several  small  pieces  of  ice  into 
a  250  cc.  bottle,  add  a  little  water,  and  about  10  gm.  of  coarse 
salt.  Fill  a  small  test  tube  half  full  of  water,  insert  a  ther- 
mometer until  the  bulb  is  immersed,  and  lower  the  test  tube 
into  the  mixture  of  ice  and  salt.  Stir  the  water  with  the 
thermometer,  and  note  the  temperature  at  which  ice  begins  to 
form  in  the  test  tube;  remove  the  test  tube,  melt  the  ice,  and 
try  again.  Make  several  trials  and  note  each  result;  take  an 
average  of  the  temperatures  observed  and  record  the  result. 

(b)  Fill  a  250  cc.  bottle  half  full  of  water,  drop  in  several 
pieces  of  ice,  and  shake  for  two  or  three  minutes.  Insert  the 
thermometer  until  the  bulb  is  immersed  and  note  the  lowest 
temperature;  repeat,  and  take  an  average,  as  in  (a). 

Experiment  16  —  Some  Chemical  Properties  of  Water 

(//  desired,  A  may  be  performed  later  as  Exp.  29  B.) 

MATERIALS.  —  Sodium,  potassium,   calcium,   test   wire,  zinc  sulphate 
solution,  sulphur,  calcium  oxide. 

A.  Interaction  with  Metals,  (a)  Sodium  and  potassium. 
(See  Precaution  in  Exp.  12  D.)  Fill  an  evaporating  dish  half 
full  of  water.  Obtain  three  or  four  small  pieces  of  sodium  from 
the  Teacher;  place  a  mortar  over  the  sodium  until  needed. 
Drop  a  piece  of  sodium  upon  the  water  in  the  dish,  stand  back 
and  observe  the  result,  waiting  for  the  slight  explosion  before 
approaching  the  dish  again ;  repeat  with  the  rest  of  the  sodium, 
piece  by  piece.  When  the  chemical  action  is  over,  stand  the 
dish  on  a  gauze-covered  ring  attached  to  an  iron  stand  (see 
Fig.  no),  and  heat  until  the  water  is  entirely  evaporated. 

Meanwhile,  proceed  with  the  potassium.  Fill  a  pneumatic 
trough  half  full  of  water.  Obtain  a  small  piece  of  potassium 
from  the  Teacher,  and  drop  the  potassium  upon  the  water. 


PROPERTIES   OF  WATER  31 

Stand  back  and  observe  the  result,  waiting  for  a  slight  explo- 
sion as  in  the  case  of  sodium. 

As  soon  as  the  water  has  been  evaporated  from  the  dish, 
examine  the  residue  as  follows:  (i)  .Moisten  the  end  of  a  glass 
rod,  touch  the  residue  with  it,  and  then  draw  this  end  across 
a  piece  of  moistened  red  litmus  paper.  Observe  the  change  in 
color  of  the  litmus  paper;  this  change  in  color  of  red  litmus 
paper  is  always  caused  by  the  strong  hydroxides  —  sodium 
hydroxide  in  this  case.  (2)  Moisten  the  looped  end  of  a 
clean  test  wire  (Fig.  103),  touch  it  to  the  residue,  and  hold 
the  end  of  the  wire  in  the  lower  and  outer  part  of  the 
Bunsen  flame.  Observe  the  color  of  the  flame;  it  is  caused 
by  the  sodium  in  the  residue.  The  production  of  this  color  is 
one  of  the  tests  for  sodium.  (3)  Dissolve  the  residue  in  5  cc. 
of  water,  pour  a  little  of  the  solution  into  a  test  tube,  add  a 
few  drops  of  zinc  sulphate  solution,  and  shake.  Observe  the 
result.  Now  pour  the  rest  of  the  solution  into  the  test  tube  and 
shake  well.  Observe  the  result.  This  is  a  test  for  the  hydroxide 
part  of  sodium  hydroxide. 

Required  Exercises. —  i.  What  property  of  water  is  shown  by 
Exp.  16  A  (a)? 

2.  State  the  chemical  change  involved  in  the  interaction  of  sodium 
and  water. 

3.  The  interaction  of  potassium  and  water  is  analogous  to  that  of 
sodium  and  water;   express  the  essential  chemical  change  in  the  form  of 
an  equation  (using  the  names  of  the  substances). 

(b)  Calcium.  Fill  a  small  test  tube  nearly  full  of  water, 
-  warm  slightly  and  stand  the  test  tube  in  the  rack.  Drop 
a  small  piece  of  clean  calcium  into  the  test  tube,  and  observe 
the  result.  If  the  action  is  not  marked,  add  another  piece  of 
calcium  or  warm  the  water.  In  a  minute  or  two,  test  the  gas 
evolved.  What  is  it?  Examine  the  contents  of  the  test  tube 
for  evidence  of  another  product.  If  the  evidence  is  doubtful, 
let  the  action  continue  and  examine  the  test  tube  subsequently, 
especially  the  lower  part. 

Recalling  the  chemical  change  in  (a),  what  chemical  change 
has  in  all  probability  taken  place  in 


32  CHEMISTRY 

B.  Combination  with  Oxides,  (a)  Put  a  little  water  in  a 
bottle.  Set  fire  to  a  small  piece  of  sulphur  in  a  deflagrating 
spoon  and  lower  the  burning  sulphur  into  the  bottle.  Let  it 
burn  a  minute  or  two,  then  extinguish  it  by  dipping  the  spoon 
into  the  water.  Remove  the  spoon,  cover  the  bottle  with  the 
hand,  and  shake  well.  Dip  a  glass  rod  into  the  liquid,  touch 
the  moistened  end  to  a  piece  of  blue  litmus  paper,  and  observe 
the  change  in  color.  This  change  in  the  color  of  blue  litmus  is 
caused  by  acids;  in  this  case  the  acid  is  sulphurous  acid,  which 
was  produced  by  the  combination  of  the  sulphur  oxide  and  water. 

(b)  Boil  a  small  piece  of  calcium  oxide  with  a  little  water  in  a 
test  tube.  Test  with  red  litmus  paper.  If  the  result  is  indiffer- 
ent, put  the  paper  in  the  test  tube  and  shake  well.  Observe 
the  result.  Recall,  or  review,  the  explanation  in  A  (a)  (i). 

Required  Exercises.  —  i.  State  briefly  the  essential  chemical  change 
that  took  place  in  Exp.  16  B  (a).  Also  in  (b). 

2.  State  these  chemical  changes  in  the  form  of  equations  (using  the 
names  of  the  substances). 

Experiment  17  —  Solubility  of  Gases 

MATERIALS.  —  As  below. 

(a)  Fill  a  test  tube  half  full  of  water,  close  with  the  thumb 
and  shake  the  test  tube  vigorously  up   and  down   several 
minutes.     Warm  the   test  tube  very  gently.     What  is  the 
immediate  evidence  of  dissolved  gas?     What  effect  has  in- 
creased heat  on  the  dissolved  gas? 

(b)  Heat  the  following  in  separate  test  tubes  as  in   (a): 
Faucet  water,  ammonia  solution,  hydrochloric  acid  solution. 
Do  the  results  resemble  those  in  (a)?     As  soon  as  the  obser- 
vation is  made,  pour  the  liquids  down  the  sink  and  flush  it  well 
with  water. 

Experiment  18  —  Solubility  of  Liquids 

MATERIALS.  —  Alcohol,    kerosene,     glycerin,     aniline,    ether,    carbon 

disulphide. 

(a)  To  a  test  tube  half  full  of  water  add  a  little  alcohol  and 
shake.  Is  there  evidence  of  solution?  Add  a  little  more  and 


PROPERTIES  OF  WATER 


33 


shake  well.  Add  a  third  portion  and  shake.  Is  there  still 
evidence  of  solution?  Draw  a  conclusion  as  to  the  solubility 
of  alcohol  in  water. 

(b)  Repeat   (a),  using  successively  kerosene,  glycerin,  ani- 
line, ether,  and  carbon  disulphide.     Observe  the  results  and 
conclude  accordingly. 

(c)  Tabulate  the  results  of  (a)  and  (b)  under  the  headings 
Mutually  Soluble,  Partly  Soluble,  Insoluble. 


Experiment  19  —  Solubility  of  Solids 

MATERIALS.  —  About  20  gm.  of  powdered  copper  sulphate,  6  gm.  of 
powdered  potassium  chlorate,  i  gm.  of  calcium  sulphate,  calcium 
hydroxide  solution,  sodium  chloride. 

A.  General,  (a)  Label  three  test  tubes,  I,  II,  III.  Put 
10  cc.  of  water  into  each.  To  I  add  i  gm.  of  powdered  copper 
sulphate,  to  II  add  i  gm.  of  powdered  potassium  chlorate,  to 
III  add  i  gm.  of  calcium  sulphate.  Shake  each  test  tube,  and 
then  allow  them  to  stand  undisturbed  for  a  few  minutes.  Is 
there  evidence  of  solubility  in  each  case?  Is  there  evidence 
of  a  varying  degree  of  solubility?  If  III  is  doubtful,  care- 
fully transfer  a  portion  of  the  clear  liquid  to  an  evaporating 
dish  by  pouring  it  down  a  glass  rod  (see  Int.  6  (i)  ),  and  evap- 
orate to  dryness.  Is  there  now  conclusive  evidence  of  solubility? 
Save  solutions  I  and  II  for  (b). 

Tabulate  the  results  of  (a)  as  follows,  using  the  customary 
terms  to  express  the  degree  of  solubility:  — 

TABLE  OF  SOLUBILITY  OF  TYPICAL  SOLIDS 


Solute 


Solvent 


Results 


1.  Copper  Sulphate 

2.  Potassium  Chlorate 

3.  Calcium  Sulphate 


Water  at  tempera- 
ture of  labora- 
tory 


34  CHEMISTRY 

(b)  Heat  I,  and  add  gradually  4  more  gm.  of  powdered  copper 
sulphate.    Does  it  all  dissolve?    Heat  II  and  add  4  more  gm. 
of  powdered  potassium  chlorate.     Does  it  all,  or  most  all, 
dissolve?      What   general   effect   has   increased   heat   on   the 
solubility  of  solids?     Save  the  solutions  for  (c). 

(c)  Heat  I  and  II  nearly  to  boiling,  and  as  the  temperature 
increases  add  the  respective  solids.     (Do  not  boil  the  solu- 
tion;   keep  it  near  the  boiling  point  by  frequent  heating.) 
Is  there  a  limit  to  their  solubility?   Draw  a  general  conclusion 
from  these  typical  results. 

B.  Special,  (a)  Fill  a  test  tube  half  full  of  clear  calcium 
hydroxide  solution,  and  heat  it  to  boiling.  Observe  the 
result.  Compare  with  the  cold  solution.  What  effect  has  heat 
upon  the  solubility  of  calcium  hydroxide? 

(b)  Prepare  a  saturated*  solution  of  sodium  chloride  by 
heating  about  5  gm.  in  10  cc.  of  water;  shake  frequently,  and 
finally  bring  the  solution  to  the  boiling  point,  but  do  not 
evaporate  much  of  the  water.  Let  the  test  tube  stand  a  minute, 
and  then  pour  the  solution  into  another  test  tube.  Observe  its 
general  appearance.  When  cool,  examine  and  compare  with 
the  clear,  hot  solution.  Answer  as  in  (a). 

Experiment  20  —  Crystallization 

MATERIALS.  —  Copper  sulphate,  alum,  potassium  dichromate,  potas- 
sium ferrocyanide,  sodium  chloride,  borax. 

Prepare  a  hot,  concentrated  solution  of  one  or  more  of  the 
following  substances,  using  about  25  cc.  of  water  and  the 
number  of  grams  indicated:  Copper  sulphate  (25),  alum  (25), 
potassium  dichromate  (15),  potassium  ferrocyanide  (25), 
sodium  chloride  (10),  borax  (5).  Pulverize  the  solid,  if  it  is 
not  provided  as  a  powder.  Prepare  the  solution  by  boiling 
the  mixture  of  water  and  the  solid  in  a  large  test  tube  several 
minutes.  Let  any  undissolved  solid  settle,  and  then  pour 
most  of  the  clear  solution  down  a  moistened  glass  rod  (see 
Int.  6  (i)  )  into  an  evaporating  dish,  or  another  small  shallow 
vessel,  taking  care  not  to  let  any  undissolved  solid  get  into 


PROPERTIES   OF   WATER  35 

the  dish.  Suspend  a  piece  of  thread  across  the  dish  and  push 
it  down  into  the  solution.  Stand  the  whole  aside  to  crystal- 
lize. Examine  at  intervals,  and  when  well-shaped  crystals 
have  formed,  especially  on  the  thread,  remove  them.  Dry 
the  crystals  carefully  with  filter  paper.  Examine  them, 
using  a  lens  if  the  crystals  are  small,  and  observe  the  proper- 
ties, particularly  the  shape,  luster,  and  color.  (Save  the 
crystals  for  later  experiments.) 

Experiment  21  —  Testing  for  Water  of  Crystallization  in 
Various  Substances 

MATERIALS.  —  See  below. 

Test  several  of  the  crystallized  substances  enumerated 
below  for  water  of  crystallization^  by  heating  a  dry  specimen 
of  each  in  a  test  tube  inclined  so  that  the  open  end  is  the 
lower:  Sodium  carbonate,  potassium  dichromate,  ferrous 
sulphate,  borax,  barium  chloride,  alum,  zinc  sulphate,  sodium 
sulphate,  calcium  sulphate,  sodium  chloride,  potassium  ni- 
trate, sugar,  magnesium  sulphate,  potassium  bromide,  and 
any  of  the  crystallized  substances  prepared  in  Exp.  20. 

Observe  in  each  case  (i)  the  change  in  appearance  of  the 
solid  during  the  heating,  (2)  relative  amount  of  water  liber- 
ated (if  appreciable),  (3)  appearance  of  the  residue. 

Experiment  22  —  Per  Cent  of  Water  of  Crystallization  in 
Crystallized  Copper  Sulphate 

MATERIALS.  —  Crystallized  copper  sulphate,  evaporating  dish. 

Clean  and  dry  an  evaporating  dish  and  weigh  it  to  a  deci- 
gram. Record  the  weight  at  once  in  the  notebook.  (See 
below.)  Powder  some  copper  sulphate  and  put  it  into  the 
dish  until  about  10  gm.  have  been  added;  then  weigh  to  a 
decigram.  Record  the  weight  at  once  (see  below).  Stand 
the  dish  with  its  contents  on  a  gauze-covered  ring  attached  to 
an  iron  stand  (see  Fig.  no)  and  heat  gently  for  five  or  ten 
minutes,  and  then  strongly  until  the  blue  color  disappears  and 


36  CHEMISTRY 

the  substance  turns  to  a  powder.  Do  not  touch  the  substance, 
and  take  special  pains  not  to  lose  any.  Cool  slowly  and 
weigh  as  before.  Record  the  weight  at  once,  and  calculate 
the  per  cent  of  water  of  crystallization.  Submit  the  result 
to  the  teacher  for  criticism. 

RECORD 

Weight  of  dish  and  substance  before  heating  =  gm. 

Weight   "     "  "  "        =  gm. 

Weight   "  "  "  "       =  gm. 

Weight  "     "       "  "         after  "       =  gm. 

Weight   "  water  of  crystallization  gm. 

Per  cent     of  water  of      "  per  cent 

Experiment  23  —  Efflorescence 

MATERIALS.  —  As  below. 

Put  a  fresh,  or  a  recently  broken,  crystal  of  several  of  the 
following  substances  on  a  piece  of  filter  paper,  and  leave  them 
exposed  to  the  air  for  an  hour  or  more:  Sodium  carbonate, 
sodium  sulphate,  borax,  ferrous  sulphate,  alum,  potassium 
ferrocyanide,  barium  chloride,  potassium  chromate,  magne- 
sium sulphate.  Describe  any  marked  change.  What  does 
the  change,  if  any,  show  about  the  air?  About  the  crystal? 
To  what  is  the  change  due? 

Experiment  24  —  Deliquescence 

MATERIALS.  —  As  below. 

Put  on  a  glass  plate  or  a  block  of  wood  a  small  piece  of 
several  of  the  following  substances:  Sodium  hydroxide,  cal- 
cium chloride,  potassium  hydroxide,  magnesium  chloride, 
table  salt,  rock  salt,  zinc  chloride,  potassium  carbonate, 
sodium  nitrate.  Leave  them  exposed  to  the  air  for  an  hour 
or  more.  Describe  any  marked  change  which  takes  place. 
What  does  the  change  show  about  the  air?  About  the  sub- 
stance? Compare  the  general  change  with  that  of  Exp.  23. 
To  what  is  the  change  due? 


PROPERTIES   OF  WATER  37 

SUPPLEMENTARY  EXPERIMENTS 
Experiment  25  —  Water  in  Various  Substances 

MATERIALS.  —  As  below  in  A. 

A.  Miscellaneous.     Proceed  as  in  Exp.  14  with  wood  (different 
kinds),  soft  coal,  fresh  grass  or  leaves,  hay,  excelsior,  raisins  or  other 
kinds  of  dried  fruit.     State  each  result. 

B.  Per  Cent  of  Water.    Devise  a  simple  experiment  to  find  the 
approximate  per  cent  of  water  in  bread,  potato,  meat,  or  some  other 
substance.     Before  proceeding,  submit  the  details  to  the  Teacher  for 
criticism. 

Experiment  26  —  Preparation  and  Properties  of  Distilled 

Water 

MATERIALS.  —  The  condenser,  etc.,  shown  in  Fig.  117,  water  contain- 
ing a  little  dirt,  calcium  chloride,  and  sodium  sulphate. 

I.  Preparation.    Fill  the  flask  C  half  full  of  the  water  containing 
the  impurities  mentioned  above,  add  a  few  short  pieces  of  glass  tubing 
to  ensure  even  boiling,  and  connect  with  the  condenser  as  shown  in 
Fig.  117.    Attach  the  inlet  (lower)  tube  to  the  faucet,  fill  the  con- 
denser slowly,  and  regulate  the  current  so  that  a  small  stream  flows 
continuously  from  the  outlet  tube  into  the  sink  or  waste  pipe.    Heat 
the  liquid  in  C  gradually,  and  when  it  boils,  regulate  the  heat  so 
that  the  boiling  is  not  too  violent.      Reject  the  first  5  or  10  cc.  of  the 
distillate,  for  they  may  contain  impurities  derived  from  the  apparatus. 
As  the  distillate  collects  in  the  clean  receiver  Z>,  proceed  as  in  II. 

II.  Properties,     (a)  Taste  a  little  distilled  water.    Compare  with 
faucet  or  well  water. 

(b)  Test  distilled  water  for  dissolved  gases  by  heating  a  little  in  a 
clean  test  tube.    State  the  result.    Compare  with  faucet  water. 

(c)  Test  distilled  water  for  organic  matter  as  follows:   Fill  a  very 
clean  test  tube  half  full  of  distilled  water,  add  a  few  drops  of  concen- 
trated sulphuric  acid,  and  enough  potassium  permanganate  solution 
to  color  the  mixture  a  light  reddish  purple.     Mix  well  by  stirring 
with  a  clean  glass  rod.    Grasp  the  test  tube  with  the  test  tube  holder 
and  heat  gently  until  the  liquid  begins  to  boil,  taking  care  to  remove 
the  test  tube  from  the  flame  occasionally  to  prevent  the  liquid  from 
spurting  out.    If  organic  matter  is  present,  the  color  of  the  solution 


38  CHEMISTRY 

will  be  changed  to  brown.    Test  in  the  same  way  some  of  the  impure 
water  used  in  I.     Compare  the  results. 

(d)  Test  separate  portions  of  distilled  water  for  mineral  matter, 
(i)  Chlorides.  Add  a  few  drops  of  nitric  acid  and  of  silver  nitrate 
solution  to  a  little  distilled  water.  Proceed  in  the  same  way  with 
some  of  the  impure  water  used  in  I.  Compare  the  results.  The  white, 
curdy  solid  is  silver  chloride,  which  is  formed  by  the  chemical  action 
between  silver  nitrate  and  the  dissolved  chloride.  All  soluble  chlorides 
produce  the  same  result.  (2)  Sulphates.  Add  a  few  drops  of  barium 
chloride  solution  to  a  little  distilled  water.  Proceed  in  the  same  way 


Fig.  117.  —  Condenser. 

with  some  of  the  impure  water  used  in  L  Compare  the  results.  The 
white,  fine  precipitate  is  barium  sulphate,  which  is  formed  by  the  chem- 
ical action  between  barium  chloride  and  the  dissolved  sulphate;  its 
formation  is  a  test  for  any  sulphate  in  solution.  (3)  Calcium  (or 
lime)  compounds.  Add  a  few  drops  of  ammonium  oxalate  solution 
to  some  distilled  water  and  also  to  some  of  the  impure  water  used  in  I. 
Compare  the  results.  The  white  precipitate  is  calcium  oxalate.  Its 
formation  serves  as  a  test  for  dissolved  calcium  compounds. 

(e)  If  time  permits,  test  samples  of  water  from  various  sources 
for  organic  and  mineral  matter. 

Experiment  27  —  Solubility  of  a  Given  Solid 

MATERIALS.  —  Solution   of   potassium   dichromate,   evaporating   dish, 

gauze-covered  ring  and  iron  stand,  water  bath. 
Weigh  an  evaporating  dish  on  the  scales,  and  record  the  weight 
in  the  notebook  (see  below).    Obtain  from  the  Teacher  about  50  cc. 


PROPERTIES   OF   WATER  39 

of  a  concentrated  solution  of  potassium  dichromate  of  known  concen- 
tration. Transfer  about  25  cc.  into  the  weighed  dish  by  a  graduate, 
noting  exactly  the  volume  taken.  (Ask  for  instructions  if  this  opera- 
tion is  not  familiar.)  Weigh  the  dish  and  contents,  and  record  the 
weight.  Stand  the  dish  on  a  water  bath  and  evaporate  to  dryness. 
While  the  solution  is  evaporating,  the  form  of  record  may  be  pre- 
pared as  shown  below;  complete  the  evaporation  by  transferring  the 
dish  to  a  gauze-covered  ring  (Fig.  1 10)  and  heating  strongly.  When 
the  dish  is  cool,  weigh,  and  record  the  weight  as  shown  below.  Heat 
again  on  the  gauze,  cool,  and  weigh;  if  the  two  weights  are  the  same 
(or  nearly  so),  accept  the  first  weighing,  but  if  the  weights  are  consider- 
ably different,  heat  intensely,  cool,  and  weigh  until  the  weight  is 
nearly  constant. 

Complete  the  entries   in  the  form  of  record,  and  calculate  the 
weight  of  the  solid  held  in  solution  by  100  gm.  of  water. 

RECORD 

(a)  Weight  of  dish    

(b)  Volume  of  solution       

(c)  Weight  of  dish  and  contents  before  heating    .... 

(d)  Weight  of  dish  and  contents  after  heating    

(e)  Weight  of  solute  (d  —  a) 

(/)  Weight  of  solvent  (c  -  d) 


Experiment  28  —  Supersaturation 

MATERIAL.  —  Sodium  thiosulphate. 

Fill  a  test  tube  half  full  of  crystallized  sodium  thiosulphate  and 
add  two  or  three  cubic  centimeters  of  water.  Warm  slowly.  As 
solution  occurs,  heat  gradually  to  boiling.  When  all  the  solid  has 
dissolved,  pour  the  solution  into  a  warm,  clean,  dry  test  tube,  insert 
a  cork  or  a  wad  of  cotton  in  the  test  tube,  and  let  it  stand  undis- 
turbed until  cool.  Observe  the  contents  and  compare  with  Exp.  20. 
Then  drop  in  a  small  crystal  of  sodium  thiosulphate  and  watch  for 
any  simple  but  definite  change.  What  happens?  Observe  and  state 
the  final  result. 


COMPOSITION    OF   WATER 


Experiment  29  —  Qualitative  Composition  of  Water 

MATERIALS.  —  As  in  Exp.  12  D  and  E  for  A,  and  in  Exp.  16  A  for  B; 
also,  for  C,  chlorine  tube  fitted  with  cork,  chlorine  water,  mortar 
or  porcelain  dish,  iron  stand  and  clamp,  joss  stick. 

A.   Hydrogen.     Recall,   perform,    or   repeat    (if   necessary) 
Exp.  12  D  and  E.      State  the  essential  result  of  each  experi- 
ment.    What   evidence    do  these  experi- 
ments give  about  the  composition  of  water? 

B.  Hydrogen  and  Oxygen.     Recall,  per- 
form, or  repeat   (if  necessary)  Exp.  16  A. 
State  the   essential    result.      What  addi- 
tional evidence  does  this  experiment  give 
about  the  composition  of  water? 

C.  Oxygen.     Obtain  250  cc.  of  chlorine 
water  from  the  Teacher.    If  fresh  chlorine 
wa"ter  is  not  available,  construct  a  chlorine 
generator,  as  described  in  Exp.  33  I,  and 
prepare  about  250  cc.  of  chlorine  water  by 
causing  the  gas  to  bubble  through  a  bottle 
of  water  until  the  water  smells  strongly 
of  the  gas. 

Fill  the  tube  with  chlorine  water,  cover 
the  open  end  with  the  thumb  or  ringer,  in- 
vert the  tube,  and  immerse  the  open  end 
in  a  mortar  or  an  evaporating  dish,  which 
should  be  nearly  full  of  chlorine  water 
(Fig.  1 1 8).  Clamp  the  tube  in  an  up- 


Fig.      1 1 8.  —  Appa- 
ratus  for  Showing 


that  Oxygen  is  a  right  position,  and  stand  the  whole  ap~ 
Constituent  of  paratus  where  it  will  receive  the  direct 
Water 

sunlight  for  several  hours.     Bubbles  of  gas 

will  soon  appear,  rise,  and  collect  at  the  top.     When  sufficient 


COMPOSITION  OF   WATER 


gas  for  a  test  has  collected,  unclamp  the  tube,  cover  the  open 
end  with  the  thumb  or  finger,  invert  the  tube,  and  put  a  glow- 
ing joss  stick  into  the  gas.  Repeat  as  long  as  any  of  the  gas 
remains.  State  the  result. 

Experiment  30  —  Electrolysis  of  Water 

MATERIALS.  —  Hofmann  apparatus,  sulphuric  acid,  joss  stick,  taper, 
matches,  platinum  tip  or  short  piece  of  capillary  glass  tubing. 
(Directions  for  making  the  platinum  tip  may  be  found  in  the 
author's  Experimental  Chemistry,  page  340.) 

Fill  the  Hofmann  apparatus  (Fig.  119)  with  water  contain- 
ing 10  per  cent  of  sulphuric  acid,  so  that  the  water  in  the 
reservoir  tube  stands  a  short  distance  above 
the  gas  tubes  after  the  stopcock  in  each 
has  been  closed.  Connect  the  platinum  ter- 
minal wires  with  a  battery  of  at  least  three 
cells.  As  the  action  proceeds,  small  bubbles 
of  gas  rise  and  collect  at  the  top  of  each 
tube.  Allow  the  current  to  operate  until  the 
smaller  volume  of  gas  is  8  to  10  cubic  centi- 
meters. Measure  the  height  of  each  gas 
column.  Assuming  that  the  tubes  have  the 
same  diameter,  the  volumes  are  in  approxi- 
mately the  same  ratio  as  their  heights. 
How  do  the  volumes  compare? 

Test  the  gases  as  follows :  (a)  Hold  a  glow- 
ing joss  stick  near  the  top  of  the  tube 
containing  the  smaller  quantity  of  gas, 
cautiously  open  the  stopcock  to  allow  the 
water  (or  air)  to  run  out  of  the  glass  tip,  Fjgf  II9  H  Hof_ 
and  then  let  out  a  little  gas  upon  the  glowing  mann  Appa- 
joss  stick.  Repeat  several  times.  What  is  ratus  for  the 
the  gas?  (b)  Open  the  other  stopcock  long  Electrolysis  of 
enough  to  force  out  the  water  (or  air)  in 
the  glass  tip;  close  the  stopcock,  and,  by  means  of  a  short 
rubber  tube,  attach  the  platinum  tip  or  the  capillary  tube 


42  CHEMISTRY 

close  to  the  end  of  the  glass  tip.  Open  the  stopcock  again, 
let  out  a  little  gas  slowly,  then  hold  a  lighted  match  for  an 
instant  at  the  end  of  the  tip,  and  immediately  thrust  a  taper 
into  the  small  and  almost  colorless  flame.  Repeat  several 
times.  What  is  the  gas? 

Required  Exercises.  —  i.  Describe  the  whole  experiment  and  sketch 
the  apparatus. 

2.  What  does  this  experiment  show  about  the  composition  of  water? 

SUPPLEMENTARY  EXPERIMENT 

Experiment  31  —  Preparation  and  Properties  of  Hydrogen 

Dioxide 

MATERIALS.  —  Three  gm.  of  barium  dioxide,  manganese  dioxide,  po- 
tassium permanganate  solution,  joss  stick,  lead  nitrate  solution, 
hydrogen  sulphide  solution. 

I.  Preparation.     Pour  about   25  cc.  of  dilute  hydrochloric  acid 
into  a  bottle  and  cool  in  running  water.    Add  slowly  about  3  gm. 
of  powdered  barium  dioxide;   stir  constantly  during  the  mixing  and 
for  several  minutes  after.    Let  the  mixture  stand  until  the  solid  settles 
somewhat,  then  filter.    If  the  filtrate  is  not  clear,  repeat  the  filtration 
through  the  same  paper  until  it  is  clear. 

II.  Properties,     (a)  Heat  a  little  of  the  filtrate  from  I;   observe 
the  result.     Now  add  a  little  powdered  manganese  dioxide  to  the 
heated  liquid  and  observe  the  result.    Test  the  escaping  gas  for  oxy- 
gen.   What  is  the  result? 

(b)  Add  several  drops  of  potassium  permanganate  solution  to  a 
little  of  the  filtrate  from  I  and  observe  the  result.     Is  a  gas  evolved? 
If  not,  add  more  potassium  permanganate  solution,  and  test  the  gas 
for  oxygen. 

(c)  Prepare  a  little  lead  sulphide  by  adding  a  few  drops  of  hydro- 
gen sulphide  solution  to  dilute  lead  nitrate  solution.    Shake  well,  add 
hydrogen  dioxide  solution  (preferably  the  commercial  solution),  and 
warm  gently.     Observe  the  result. 

(d)  Examine  the  inner  end  of  the  cork  stopper  of  a  bottle  of  hydro- 
gen dioxide.    Explain. 


LAW  OF  CONSTANT  COMPOSITION 

SUPPLEMENTARY  EXPERIMENT 

(See  note  to  Exp.  13.) 

Experiment  32  —  The  Combination  of  Oxygen  with 
Magnesium 

MATERIALS.  —  Porcelain   crucible   and   cover,    powdered   magnesium, 
forceps,  pronged  tripod  (or  iron  ring  and  triangle),  crucible  block. 

Clean  and  dry  the  crucible  and  cover,  and  weigh  both  together 
on  the  balance.  Record  the  weight  in  the  notebook  as  shown  below. 
Put  from  .4  to  .5  gm.  of  magnesium  in  the  crucible,  and  weigh  again 
(including  cover).  Record  the  weights  thus:  — 

Weight  of  crucible,  cover,  and  magnesium  ..... 

Weight  of  crucible  and  cover    ................ 

Weight  of  magnesium      ..................... 

Stand  the  crucible  on  the  tripod,  as  shown  in  Fig.  120  (or  on  a 
triangle  supported  by  a  ring),  and  heat  for  five 
minutes  with  a  flame  which  just  touches  the  bot- 
tom of  the  crucible.  Grasp  the  cover  firmly  by  the 
ring  with  the  clean  forceps,  cautiously  lift  it,  and 
if  the  magnesium  glows,  cover  the  crucible  in- 
stantly. Repeat  this  operation  at  frequent  in- 
tervals, gradually  increasing  the  heat,  until 

the    glow   ceases  to  spread 

^          ,      ,  ,  ,     rig.  120.  —  Covered 

through  the  mass;   then  ad- 


just  the  cover  so  that  a  small       p()rted  Qn  a  Tri. 

opening  is  left  between  the        p0(j 

cover  and  the  crucible,  and 

heat  strongly  for  ten  or  fifteen  minutes.     If  the 

contents  has  ceased  to  glow,  heat  the  crucible, 
Fig.  121.  —  Crucible  '.  . 

,3,     ,   ,      ~  uncovered,  for  five  or  ten  minutes.     Take  care 

Block  for  Carry- 

ing a  Crucible        not  to  uPset  tne  cover  by  accident  or  insecure 

handling  with  the  forceps.     At  no  time,  should 

the  flame  touch  the  cover  of  the  crucible;  generally  speaking,  the 

flame  should  reach  as  high  outside  as  the  magnesium  does  inside. 


44  CHEMISTRY 

Cool  the  crucible  gradually.  When  cool  enough  to  touch,  it  is  cool 
enough  to  weigh.  In  carrying  the  crucible  to  and  from  the  balance, 
it  should  be  placed  in  the  crucible  block  (Fig.  121).  Weigh  and 
record  in  the  notebook  thus:  — 

Weight  of  crucible,  cover,  and  contents,  after  heating 

Weight  of  crucible,  cover,  and  contents,  before  heating 

Weight  of  oxygen  which  has  combined  with  the  magnesium  . . 

Heat  the  uncovered  crucible  again  strongly  for  five  or  ten  mintues, 
cool,  and  weigh  as  before.  If  the  weight  is  not  the  same,  continue 
until  the  last  two  weights  are  approximately  equal.  Record  each 
weight. 

From  the  weights  of  the  magnesium  taken  and  the  oxygen  found, 
calculate  the  ratio  in  which  the  two  elements  combined.  Submit  the 
result  to  the  Teacher  for  criticism  before  throwing  away  the  contents 
of  the  crucible. 

NOTE. — The  crucible,  if  blackened,  can  be  cleaned  by  heating  a  little 
sodium  hydroxide  in  it  and  then  washing  thoroughly  with  water. 


CHLORINE  — ACIDS,   BASES,   SALTS 


Experiment  33  —  Preparation  and  Properties  of  Chlorine 

(Perform  this  experiment  in  the  hood.} 

MATERIALS.  —  Concentrated  hydrochloric  acid,  30  gm.  of  manganese 
dioxide,  wad  of  iron  thread,  cotton,  calico,  paper  with  writing  in 
lead  pencil  and  in  ink,  litmus  paper  (both  colors),  taper,  and  a 
piece  of  copper  wire  15  cm.  long.  The  apparatus  is  shown  in 
Fig.  122.  A  is  a  250  cc.  Erlenmeyer  flask  which  stands  on  a 
gauze-covered  ring;  the  parts  lettered  B,  C,  D,  E  have  been  used 
in  preceding  experiments.  There  are  also  needed  four  bottles 
like  G,  a  wooden  block  F  (about  10  cm.  or  4  in.  square)  with  a  hole 
in  the  center,  four  glass  plates  to  cover  the  bottles. 

I.  Preparation.  Weigh  the  manganese  dioxide  upon  a 
piece  of  paper  creased  lengthwise,  and  slip  it  into  the  flask. 
Arrange  the  apparatus  as  shown 
in  Fig.  122.  Introduce  enough 
concentrated  hydrochloric  acid 
through  the  dropping  tube  B  to 
cover  the  manganese  dioxide.  Heat 
the  flask  A  gently  with  a  small 
flame.  Chlorine  is  evolved  as  a 
greenish  yellow  gas,  and  passes 
into  the  bottle  G,  which  should  be 
removed  when  full  (as  seen  by  the 
color)  and  covered  with  a  glass 
plate;  the  bottle  may  be  easily  re- 
moved by  holding  the  block  F  in 
one  hand  and  pulling  the  bottle  G 
aside,  bending  the  whole  delivery 
tube  at  the  same  time  at  the  rub- 
ber connection  D.  If  the  evolution 
of  gas  slackens,  introduce  more  acid, 
and  proceed  at  once  with  II. 


Fig.  122.  —  Apparatus  for 
Preparing  Chlorine. 

Collect  four   bottles, 


46  CHEMISTRY 

II.  Properties,  (a)  Twist  one  end  of  the  copper  wire 
around  a  wad  of  iron  thread  (Fig.  123),  heat  the  edge  of  the 
wad  for  an  instant  in  the  flame,  and  quickly  lower  it  into  a 
bottle  of  chlorine.  Observe  the  result.  Dissolve  the  contents 
of  the  bottle  in  a  little  water,  filter  if  not  clear,  and  test  the 
clear  solution  for  a  chloride  (see  Exp.  26  II  (d)  (i)).  State 
the  result. 

(b)  Into  a  bottle  of  dry  chlorine  put  a  piece  of  calico,  litmus 
paper  (both  colors),  and  paper  containing  writing  in  black 
and  in  red  ink.  Allow  the  whole  to  remain 
undisturbed  for  a  few  minutes  and  then  ob- 
serve the  change  in  the  materials,  if  any. 
Add  several  drops  of  water,  shake  the 
bottle,  and  then  observe  the  change.  Draw 
a  general  conclusion  from  the  whole  experi- 
ment. 

(c)  Lower  a  burning  taper  a  short  distance 
into   a  bottle  of  chlorine,  and  observe  the 
two  products   as  the  taper  burns.    Draw  a 
conclusion.      Verify     it    thus :     Twist     the 
Fig.  123.  —  Wads  other   end  of  the   copper   wire   used  in  (a) 
of  Cotton  and   arOund  a  piece  of  cotton   (Fig.    123);    cau- 
Iron  Thread.       tiously  h^   about    10  cubic  centimeters  of 
turpentine  in  a  large  test  tube, *  saturate  the  cotton  with  the 
hot  turpentine,  and  lower  the  cotton  into  a  bottle  of  chlo- 
rine.    Observe  the  result,  especially  at  the  beginning  of  the 
reaction. 

NOTE.  —  As  soon  as  II  (c)  is  performed,  pour  the  contents  of  the 
flask  into  a  waste  jar  in  the  hood.  The  bottle  used  in  the  latter  part  of 
(c)  may  be  cleaned  by  adding  water,  a  little  sand,  and  several  pieces  of 
paper,  and  then  shaking  vigorously. 

1  Hold  the  test  tube  with  the  holder.  Remember  that  turpentine 
ignites  easily.  If  the  turpentine  catches  fire,  press  a  damp  towel  over 
the  mouth  of  the  test  tube. 


CHLORINE  —  ACIDS,   BASES,   SALTS  47 

Experiment  34  —  The  Characteristic  Property  of  Bleaching 

Powder 

MATERIALS.  —  Bleaching  powder,  dilute  sulphuric  acid,  colored  cloth, 
unbleached  cloth,  glass  rod,  evaporating  dish,  two  bottles. 

Put  a  little  bleaching  powder  into  an  evaporating  dish,  and 
add  enough  water  to  make  a  thin  paste.  Add  10  cubic  centi- 
meters of  dilute  sulphuric  acid  to  a  bottle  half  full  of  water. 
Fill  the  other  bottle  nearly  full  of  water.  Tear  off  a  small 
piece  of  the  colored  cloth  for  a  sample.  Dip  the  rest  of  the 
colored  cloth  into  the  bleaching  powder  and  then  into  the 
acid,  passing  it  back  and  forth  several  times.  Finally  wash 
the  cloth  thoroughly  in  the  bottle  of  water,  squeeze  out  the 
excess  of  water,  let  the  washed  cloth  dry  somewhat,  and  then 
compare  its  color  with  that  of  the  sample.  Describe  the 
change  in  the  appearance  of  the  cloth. 

Proceed  in  the  same  way  with  the  unbleached  cloth,  and 
state  the  result. 

Experiment  35  —  Preparation  and  Properties  of  Hydrogen 
Chloride  and  Hydrochloric  Acid 

(Perform  this  experiment  in  the  hood.) 

MATERIALS.  —  The  apparatus  shown  in  Fig.  122;  20  grams  of  sodium 
chloride,  concentrated  sulphuric  acid,  joss  stick,  litmus  paper 
(blue),  ammonium  hydroxide. 

I.  Preparation,  (a)  Hydrogen  chloride.  Put  8  cubic 
centimeters  of  water  into  a  small  bottle  or  an  evaporating 
dish,  cautiously  add  12  cubic  centimeters  of  concentrated 
sulphuric  acid,  and  stir  until  the  two  are  mixed.  While  this 
mixture  is  cooling,  weigh  the  sodium  chloride,  slip  it  into  the 
flask,  and  arrange  the  apparatus  as  shown  in  Fig.  122.  Intro- 
duce half  the  cold  acid  mixture  through  the  tube,  let  it  settle 
through  the  sodium  chloride,  and  then  introduce  the  remain- 
ing acid.  Heat  the  flask  gently  with  a  low  flame,  as  in  the 
preparation  of  chlorine.  Hydrogen  chloride  is  evolved,  and 
passes  into  the  bottle,  which  should  be  removed  when  full,  as 
directed  under  chlorine.  A  piece  of  moist  blue  litmus  paper 


48  CHEMISTRY 

held  at  the  mouth  of  the  bottle  will  show  when  it  is  full.  Col- 
lect three  bottles  of  the  gas,  cover  each,  when  filled,  with  a 
glass  plate,  and  set  aside  until  needed  for  II. 

(b)  Hydrochloric  acid.  As  soon  as  the  third  bottle  of  gas 
has  been  collected,  removed,  and  covered,  put  in  its  place  a 
bottle  one  fourth  full  of  water.  Adjust  the  delivery  tube  E 
so  that  the  lower  end  is  a  short  distance  above  the  surface  of 
the  water.  Continue  to  heat  the  flask  at  intervals,  and  the 
gas  will  be  absorbed  by  the  water.  Shake  the  bottle  occasion- 
ally. Meanwhile  perform  II. 

II.  Properties    of   hydrogen   chloride.     Proceed    as    follows 
with  the  gas  prepared  in  I  (a):  —  (a)  Insert  a  blazing  joss 
stick  once  or  twice  into  one  bottle,  and  observe  the  result. 
Compare   the   behavior   of  hydrogen   chloride   with  that  of 
hydrogen  and  oxygen  under  similar  conditions. 

(b)  Hold  a  piece  of  wet  filter  paper  near  the  mouth  of  the 
same  bottle.     Observe  and  describe  the  result.     What  is  the 
cause? 

(c)  Invert  a  bottle  of  the  gas,  and  stand  it  in  a  vessel  of  water 
(e.g.  the  pneumatic  trough).     Observe  any  change  inside  the 
bottle  after  a  few  minutes.     What  property  of  the  gas  does 
the  result  illustrate?     Verify  the   observation  by  a  simple 
test  applied  to  the  contents  of  the  bottle. 

(d)  Drop  into  the  remaining  bottle  of  gas  a  piece  of  filter 
paper  wet  with  ammonium  hydroxide.     Describe  the  result. 
What  is  the  name  of  the  product? 

(e)  State  other  properties  of  hydrogen  chloride  which  you 
have  observed,  e.g.  'color,  odor,  density,  behavior  with  litmus. 

III.  Properties  of  hydrochloric  acid.     Remove  the  bottle  in 
which  the  hydrogen  chloride  is  being  absorbed,  and  study  the 
aqueous  solution  of  the  gas  as  follows :  —  (a)  Determine  its 
general  properties,  e.g.  taste  (cautiously),  action  with  litmus, 
and  with  magnesium.     State  the  results. 

(b)  Add  to  a  test  tube  half  full  of  the  hydrochloric  acid  a  few 
drops  of  nitric  acid  and  of  silver  nitrate  solution.  Describe 
the  precipitate.  What  is  its  name?  Shake  the  test  tube 


CHLORINE  —  ACIDS,   BASES,   SALTS  49 

filter  part  of  the  contents,  and  expose  the  precipitate  upon  the 
paper  to  the  sunlight.  Describe  the  change  in  the  precipitate 
which  soon  occurs.  To  the  remaining  contents  of  the  test 
tube  add  considerable  ammonium  hydroxide,  and  shake. 
Describe  the  result. 

NOTE.  —  As  soon  as  III  (b)  is  performed,  add  water  to  the  flask, 
shake  well,  and  pour  the  contents  into  a  waste  jar  in  the  hood. 

Experiment  36  —  Tests  for  Hydrogen  Chloride, 
Hydrochloric  Acid,  and  Chlorides 

(a)  Recall  properties  which  would  serve  as  a  test  for  hydrogen 
chloride. 

(b)  Apply  (a)  to  hydrochloric  acid. 

(c)  Suggest  a  test  for  a  soluble  chloride.     Apply  it  to  several 
chlorides,  especially  some  not  used  in  previous  experiments. 

Experiment  37  —  General  Properties  of  Acids 

MATERIALS.  —  Dilute  sulphuric,  nitric,  and  hydrochloric  acids,  glass 
rod,  litmus  paper  (both  colors),  magnesium. 

Fill  three  test  tubes  one-third  full  of  water;  add  a  few  drops 
of  concentrated  sulphuric  acid  to  one,  of  concentrated  hydro- 
chloric acid  to  another,  and  of  concentrated  nitric  acid  to  the 
third.  Shake  each  test  tube  thoroughly,  and  label  them  in 
some  distinguishing  manner.  Determine  the  general  proper- 
ties of  the  acids  as  follows:  — 

(a)  Dip  a  clean  glass  rod  into  each  acid  successively  and 
cautiously  taste  it.     Describe  the  taste  by  a  single  word. 

(b)  Dip  a  clean  glass  rod  into  each  acid  successively  and 
put  a  drop  on  both  kinds  of  litmus  paper.     Describe  the 
change.     The  striking  change  is  characteristic  of  acids. 

(c)  Slip  a  small  piece  of  magnesium  into  the  test  tubes 
containing    the    sulphuric    and    hydrochloric    acids.     If    no 
chemical  action  results,  warm  gently.     Test  the  most  obvious 
product  by  holding  a  lighted  match  inside  of  each  tube.      What 
gas  comes  from  the  hydrochloric  and  sulphuric  acids? 

(e)  Summarize  the  general  results  of  this  experiment. 


50  CHEMISTRY 

Experiment  38  —  General  Properties  of  Bases 

MATERIALS.  —  Sodium  hydroxide  and  potassium  hydroxide  solutions, 
ammonium  hydroxide,  litmus  paper  (both  colors),  glass  rod. 

Determine  the  general  properties  of  bases  as  follows:  - 

(a)  Rub  a  little  of  each  solution  between  the  fingers,  and 
describe  the  feeling. 

(b)  Cautiously  taste  each  liquid  by  touching  to  the  tip  of  the 
tongue  a  rod  moistened  with  each,  and  describe  the  result. 

(c)  Test  each  solution  with  litmus  paper.     Describe  the 
result. 

(d)  Summarize  the  general  results  of  this  experiment.    Com- 
pare acids  and  bases  as  to  taste  and  to  reaction  with  litmus. 

Experiment  39  —  A  Property  of  Many  Salts 

MATERIALS.  —  Litmus  paper  (both  colors),  glass  rod,  dilute  solutions 
of  chemically  pure  sodium  chloride,  potassium  nitrate,  potassium 
sulphate,  barium  chloride,  potassium  chlorate,  potassium  bromide, 
and  strontium  nitrate. 

Test  the  solutions  with  litmus  paper.  Describe  the  result 
in  each  case.  Compare  the  litmus  reaction  of  these  salts  with 
the  reaction  of  acids  and  bases. 

Experiment  40  —  Neutralization 

MATERIALS.  —  Sodium  hydroxide  (solid),  hydrochloric  acid,  blue 
litmus  paper,  glass  rod,  evaporating  dish,  gauze-covered  ring. 

Dissolve  a  small  piece  of  sodium  hydroxide  in  an  evap- 
orating dish  one-third  full  of  water.  Add  a  little  dilute 
hydrochloric  acid,  and  stir  well;  continue  to  add  the  acid, 
until  a  drop  of  the  well  mixed  solution  taken  from  the  dish 
upon  a  clean  glass  rod  just  reddens  blue  litmus  paper.  Then 
evaporate  the  solution  to  dryness  by  heating  the  dish  on  a 
gauze-covered  ring  (Fig.  no).  Since  the  residue  retains 
traces  of  the  excess  of  hydrochloric  acid  added,  it  is  necessary 
to  evaporate  all  of  this  acid  before  applying  any  test.  Heat 
the  dish  until  the  yellow  color  disappears,  then  moisten  the 


CHLORINE  —  ACIDS,  BASES,  SALTS     51 

whole  residue  carefully  with  a  litte  warm  water,  and  heat 
again  to  evaporate  the  last  traces  of  acid;  it  is  advisable  to 
add  and  evaporate  two  portions  of  water. 

Test  a  portion  of  the  residue  with  moist  litmus  paper  to 
find  whether  it  has  acid,  basic,  or  neutral  properties.  Taste 
a  little.  Test  (a)  a  solution  of  a  little  of  the  residue  for  a 
chloride,  and  (b)  a  portion  of  the  solid  residue  for  sodium  by 
heating  a  little  on  a  test  wire  in  the  flame.  What  is  the 
residue? 

SUPPLEMENTARY  EXPERIMENTS 

Experiment  41  —  Preparation  of  Chlorine  from  Various 
Substances 

(Each  pupil  need  not  perform  all  of  this  experiment.) 

MATERIALS.  —  Sodium  chloride,  manganese  dioxide,  concentrated 
hydrochloric  acid,  potassium  chlorate,  potassium  permanganate, 
potassium  dichromate,  lead  tetroxide,  lead  dioxide. 

A.  Put  a  little  sodium  chloride  and  manganese  dioxide  in  a  test 
tube,  mix  thoroughly  by  shaking,  add  a  little  dilute  sulphuric  acid, 
and  warm  gently.      Observe  the  color  of  the  liberated  gas,  its  odor 
(very  cautiously),  and  its  action  upon  moist  litmus  paper.      What 
is  the  gas? 

B.  Put  a  few  crystals  of  potassium  chlorate  in  a  test  tube,  add  a 
little  dilute  hydrochloric  acid,  and  warm  gently.     Observe  and  test 
the  gaseous  product  as  in  A.     What  is  the  gas? 

C.  Put  a  few  crystals  of  potassium  permanganate  in  a  test  tube, 
add  not  more  than   5  cc.  of  concentrated  hydrochloric  acid,  and 
observe  and  test  the  gaseous  product  as  in  A.     What  is  the  gas? 

D.  Proceed  as  in  C,  using  potassium  dichromate.     Warm  gently. 
What  gas  is  produced? 

E.  (a)  Proceed  as  in  D,  using  lead  tetroxide  (red  lead).      What 
gas  is  produced?     (6)  Repeat  (a),  using  lead  dioxide. 


52  CHEMISTRY 

Experiment  42. —  Preparation  of  Hydrogen  Chloride  from 
Various  Substances 

(Each  pupil  need  not  perform  all  of  this  experiment.) 

MATERIALS.  —  Concentrated  hydrochloric  acid,  concentrated  sulphuric 
acid,  silver  nitrate  solution,  litmus  paper,  ammonium  chloride, 
barium  chloride,  calcium  chloride. 

A.  Put  a  little  concentrated  hydrochloric  acid  into  a  test  tube, 
add  a  few  drops  of  concentrated  sulphuric  acid,  and  test  the  escaping 
gas  with  (a)  moist  blue  litmus  paper,  (b)  moist  filter  paper,  and  (c) 
a  glass  rod  to  the  end  of  which  a  little  silver  nitrate  solution  adheres. 
State  the  results.    What  is  the  gas. 

B.  Put  a  little  ammonium  chloride  into  a  test  tube,  add  a  few 
drops  of  concentrated  sulphuric  acid,  warm  slightly,  if  necessary, 
and  test  the  escaping  gas  as  in  A.     State  the  results.     What  is  the 
gas? 

C.  Repeat  B,  using  several  chlorides,  e.g.  barium  chloride,  calcium 
chloride.     State  the  results.     Draw  a  general  conclusion. 

Experiment  43  —  Aqua  Regia 

MATERIALS.  —  Gold  leaf,  concentrated  nitric  and  hydrochloric  acids, 
glass  rod. 

Touch  a  small  piece  of  gold  leaf  with  the  end  of  a  moist  glass  rod, 
and  wash  the  gold  leaf  into  a  test  tube  by  pouring  a  few  cubic  centi- 
meters of  concentrated  hydrochloric  acid  down  the  rod.  Warm  gently. 
Does  the  gold  dissolve?  Wash  another  piece  of  gold  leaf  from  a 
clean  glass  rod  very  carefully  into  another  test  tube  with  concen- 
trated nitric  acid.  Heat  as  before.  Does  the  gold  dissolve?  Pour 
the  contents  of  one  tube  cautiously  into  the  other.  Warm  gently, 
if  no  change  occurs.  Does  the  gold  dissolve? 

Required  Exercises.  —  i.  What  compound  of  gold  is  formed  by  its 
interaction  with  aqua  regia? 

2.  Would  chlorine  water  act  like  aqua  regia  upon  gold?  (If  in  doubt, 
try  the  experiment.) 


CHLORINE  —  ACIDS,   BASES,   SALTS  53 

Experiment  44  —  Litmus  Reaction  of  Some  Common 
Substances 

MATERIALS.  —  "Lemon  juice,  vinegar,  sweet  and  sour  milk,  washing 
soda,  borax,  wood  ashes,  faucet  water,  baking  soda,  sugar,  cream 
of  tartar,  alum,  soap,  tooth-powder,  the  juice  of  any  ripe  fruit  and 
any  unripe  fruit,  household  ammonia,  potash,  limewater,  pickles, 
jelly,  grape  juice. 

Apply  the  litmus  test  to  the  substances  enumerated  above.  Make 
a  solution  of  each  of  the  solids  before  testing.  Tabulate  the  results 
under  the  terms,  Acid,  Basic,  Neutral. 


NITROGEN  — NITROGEN   COMPOUNDS 


Experiment  45  —  Preparation  and  Properties    of  Nitrogen 

MATERIALS.  —  Apparatus  as  shown  in  Fig.  124,  three  bottles,  joss 
stick,  iron  thread,  small  piece  of  sulphur  and  a  deflagrating  spoon, 
10  grams  of  ammonium  chloride  and  10  grams  of  sodium  nitrite. 

I.  Preparation.      Weigh  the  two    substances,  put  them  in 
the  flask,  and  add  50  cc.  of  water.      Arrange  the  apparatus  as 

shown  in  Fig.  124.  Fill 
the  cup  of  the  dropping 
funnel  with  water,  and  then 
ask  to  have  the  apparatus 
inspected. 

Heat  the  flask  gently 
with  a  low  flame  and  as 
soon  as  the  nitrogen  bub- 
bles regularly  through  the 
water,  slip  the  collecting 
bottle  over  the  hole  in  the 
support.  Heat  gently,  but 
enough  to  keep  the  gas 
bubbling  slowly  through 
Fig.  124.  — Apparatus  for  Preparing  the  water  Collect  three 

bottles  of  nitrogen.  Cau- 
tion. —  If  the  mixture  in  the  flask  begins  to  froth  or  the  gas 
comes  off  too  rapidly,  remove  the  flame  and  let  in  a  little 
water;  if  it  continues  to  froth,  pinch  the  clamp  and  let  out 
the  excess  of  gas.  As  soon  as  the  frothing  ceases,  close  the 
clamp  and  continue  to  heat.  Remove  the  delivery  tube  as 
soon  as  the  three  bottles  of  nitrogen  have  been  collected. 
Proceed  at  once  with  II. 

II.  Properties,     (a)  Thrust    a    blazing    joss    stick    into    a 
bottle  of  the  gas.    Observe  and  state  the  result. 

(b)  Put  a  small  piece  of  sulphur  in  a  deflagrating  spoon, 
light  the  sulphur,  lower  it  into  a  bottle  of  nitrogen,  and  keep 


NITROGEN  — NITROGEN    COMPOUNDS         55 

/ 
it  there  about  half  a  minute.    Observe  the  result.    Withdraw, 

and  observe  the  result.    State  the  results. 

(c)  Wind  one  end  of  a  copper  wire  around  a  wad  of  iron 
thread,  kindle  it  along  one  edge,  and  quickly  thrust  the  glow- 
ing iron  into  a  bottle  of  nitrogen.  Observe  and  state  the 
result. 

Required  Exercises.  —  i.  Describe  briefly  the  preparation  of 
nitrogen. 

2.  Sketch  the  apparatus,  if  time  permits. 

3.  Compare  the  characteristic  properties  of  nitrogen  with  those  of 
oxygen  found  by  similar  experiments. 

Experiment  46  —  Preparation  and  Properties  of  Ammonia 
Gas  and  Ammonium  Hydroxide 

(Perform  this  experiment  in  the  hood.) 

MATERIALS.  —  15  grams  of  lime  (calcium  oxide),  15  grams  of  am- 
monium chloride,  3  bottles,  2  glass  plates,  pneumatic  trough  filled 
as  usual,  litmus  paper,  joss  stick,  filter  paper.  The  apparatus  is 
shown  (in  part)  in  Fig.  125.  The  flask  A  is  provided  with  a  one- 
hole  rubber  stopper  to  which  is  fitted  the  right-angle  bend  C 
connected  with  a  glass  tube  B  (12  centimeters  or  5  inches  long) 
by  the  rubber  tube  D. 

I.  Preparation,  (a)  Ammonia  gas.  Weigh  the  lime 
and  ammonium  chloride  separately,  and  mix  them  thoroughly 
on  a  piece  of  paper.  Slip  the  mixture  into  the  flask,  and 
add  a  little  water,  thereby  transforming  the  calcium  oxide 
into  calcium  hydroxide.  Quickly  insert  the  stopper  with  its 
tubes,  and  clamp  the  flask  as  shown  in  Fig.  125. 

Slip  the  glass  delivery  tube  B  into  a  bottle,  invert  the  bottle, 
and  hold  it  so  that  the  tube  is  in  the  position  shown  in  the 
figure.  Heat  the  flask  gently  with  a  low  flame.  Ammonia 
gas  will  pass  up  into  the  bottle,  which  should  be  removed, 
when  full,  and  covered  with  a  glass  plate.  A  piece  of  moist 
red  litmus  paper  held  near  the  mouth  will  show  (by  change  in 
color)  when  the  bottle  is  full.  Do  not  smell  at  the  mouth  of 
the  bottle.  Collect  two  bottles  and  set  them  aside  until  needed 
for  II. 


CHEMISTRY 


(b)  Ammonium  hydroxide.  As  soon  as  the  last  bottle  has 
been  collected,  rearrange  the  apparatus  to  absorb  the  am- 
monia gas  in  water,  as  in  the  case  of  hydrochloric  acid  (see 
Exp.  35  I  (b)).  Replace  the  glass  tube  B  by  the  delivery 
tube  E,  which  should  pass  through  the  wooden  block  F  into 
a  bottle  G  one-fourth  full  of  water,  so  that  the  end  is  just 
above  the  surface  of  the  water.  Continue  to  heat  the  flask 

gently  at  intervals,  and  the 
gas  will  be  absorbed  by  the 
water.  Shake  the  bottle  oc- 
casionally. Meanwhile  per- 
form II. 

II.  Properties  of  ammonia 
gas.  Proceed  as  follows  with 
the  ammonia  gas  prepared  in 
I  (a) :  —  (a)  Test  the  gas  in 
one  bottle  with  a  blazing  joss 
stick.  Observe  the  result. 
Compare  the  behavior  of  am- 
monia gas  with  that  of  hy- 
drogen, oxygen,  and  hydrogen 


Fig.  125.  —  Apparatus  for  Prepar- 
ing Ammonia. 


chloride  under  similar  circum- 
stances. 

(b)  Invert  the  same  bottle 


in  the  pneumatic  trough,  and  shake  it  vigorously,  taking 
care  to  keep  the  mouth  under  water.  Observe  any  change 
noticed  inside  the  bottle  after  a  few  minutes.  What  property 
of  the  gas  is  revealed?  Is  it  a  marked  property?  Test  the 
contents  of  the  bottle  with  litmus  paper  (both  colors),  and 
state  the  result. 

(c)  Pour  a  few  drops  of  concentrated  hydrochloric  acid  into 
an  empty,  warm,  dry  bottle.  Rotate  the  bottle  until  the  inside 
is  well  moistened  with  the  acid.  Cover  it  with  a  glass  plate, 
invert  it,  and  stand  it  upon  a  covered  bottle  of  ammonia  gas. 
Remove  both  plates  at  once,  and  hold  the  bottles  together  by 
grasping  them  firmly  about  their  necks.  Observe  the  result. 


NITROGEN— NITROGEN   COMPOUNDS          57 

Describe  the  result,  giving  the  evidence  of  the  chemical  action. 
What  is  the  white  substance? 

(d)  State  other  properties  of  ammonia  gas  you  have  ob- 
served, e.g.  color,  odor,  density,  and  behavior  with  litmus 
paper. 

III.  Properties  of  ammonium  hydroxide.  Remove  the  bottle 
in  which  the  ammonia  gas  is  being  absorbed  in  I  (b),  and 
proceed  with  the  resulting  ammonium  hydroxide  as  follows :  — 

(a)  Determine  the  general  properties,  e.g.  taste  and  odor 
(cautiously),  feeling,  behavior  with  litmus  paper. 

(b)  Warm  a  little  in  a  test  tube.     What  gas  is  evolved? 
Continue  the  heating,  and  test  the  escaping  gas  frequently  by 
the  odor  (cautiously) ;   state  the  result. 

(c)  Put  a  few  cubic  centimeters  of  the  ammonium  hydroxide 
in  an  evaporating  dish,  stand  the  dish  in  the  hood  or  in  the 
open  air,  and  in  an  hour  (or  before  the  liquid  evaporates  com- 
pletely) test  the  solution  by  the  odor.    State  the  final  result. 

NOTE.  —  As  soon  as  the  bottle  of  ammonium  hydroxide  is  removed 
from  E  (in  the  generating  apparatus)  the  stopper  of  the  flask  should  be 
loosened;  subsequently  the  contents  of  the  flask  should  be  thrown 
into  a  waste  jar  in  the  hood. 

Experiment  47  —  Preparation  of  Nitric  Acid 

Precaution.  Nitric  acid  is  very  corrosive,  and  may  cause  a  serious 
burn  if  it  comes  in  contact  with  the  skin. 

MATERIALS.  —  Glass  stoppered  retort,  iron  stand,  ring,  gauze,  bottle, 
30  grams  of  sodium  nitrate,  20  cubic  centimeters  of  concentrated 
sulphuric  acid,  funnel. 

Weigh  the  sodium  nitrate  and  slip  it  into  the  retort  through 
the  tubulure.  Fill  the  bottle  nearly  full  of  water.  Put  a  large 
empty  test  tube  into  the  bottle,  insert  the  neck  of  the  retort 
into  the  test  tube,  and  clamp  the  apparatus  as  shown  in  Fig. 
126.  Stand  a  funnel  in  the  tubulure  of  the  retort  so  that  the 
end  is  well  inside  the  bulb,  and  pour  the  acid  very  carefully 
through  the  funnel.  Remove  the  funnel  and  insert  the  stopper 
of  the  retort  tightly.  Heat  the  retort  gently  as  long  as  any 


CHEMISTRY 


nitric  acid  runs  down  the  neck  into  the  test  tube.  Then  un- 
clamp  the  retort,  and  remove  the  test  tube  carefully.  Leave 
the  nitric  acid  in  the  test  tube  until  needed  for  Exp.  48,  cork- 
ing the  test  tube  unless  the  acid  is  to  be  used  soon. 

NOTE.  —  Allow  the  contents  of  the  retort  to  cool,  add  a  little  water, 
boil  until  the  solid  in  the  bulb  is  reduced  to  a  small  bulk  or  dissolved, 
and  pour  it  into  a  waste  jar  in  the  hood. 

Experiment  48  —  Some  Properties  of  Nitric  Acid 

MATERIALS.  —  Concentrated  and  dilute   nitric  acid,   quill   toothpick, 
sulphur,  zinc,  magnesium. 

A.  Concentrated,  (a)  Observe  the  color  of  the  concentrated 
nitric  acid  prepared  in  Exp.  47.  Compare  it  with  the  con- 
centrated nitric  acid  in 
several  bottles  in  the 
laboratory  and  with  the 
typical  specimen  of  con- 
centrated nitric  acid 
placed  upon  the  side 
shelf  by  the  Teacher. 
State  the  result. 

(b)  Hold  a  piece  of  wet 
filter  paper  at  the  mouth 
of  the  test  tube  (or  a 
bottle)  of  concentrated 

Fig.  126.  —  Apparatus  for  Preparing       nitric  acid-     Observe  and 
Nitric  Acid.  state   the  result.      Com- 

pare    the     behavior     of 
nitric  acid  with  that  of  concentrated  hydrochloric  acid. 

(c)  Repeat   (b),   using   a  piece   of  filter  paper   moistened 
with  ammonium  hydroxide.    What  is  the  name  of  the  product? 

(d)  Pour  5  cubic  centimeters  of  concentrated  nitric  acid 
very  carefully  into  a  test  tube,  drop  in  a  piece  of  a  quill  tooth- 
pick, and  observe  any  change  in  the  color  of  the  quill.    Heat 
very  gently,  and  observe  the  effect  upon  the  quill.    State  the 
final  result. 


NITROGEN— NITROGEN    COMPOUNDS          59 

(e)  Put  about  i  gram  of  sulphur  in  a  test  tube,  add  care- 
fully 5  cubic  centimeters  of  concentrated  nitric  acid,  attach 
the  test  tube  holder,  and  boil  very  cautiously  —  in  the  hood  — 
for  a  few  minutes.  Add  10  to  15  cubic  centimeters  of  water, 
filter  the  solution,  if  it  is  not  clear,  and  test  the  filtrate  for  a 
sulphate  by  adding  barium  chloride  solution.  State  the 
result. 

(/)  Stand  three  test  tubes  in  the  test  tube  rack,  put  a 
piece  of  zinc  into  one,  copper  into  another,  and  magnesium 
ribbon  (rolled  into  a  ball)  into  the  third.  Add  a  little  concen- 
trated nitric  acid  to  each  test  tube.  Observe  the  result.  Test 
the  gaseous  product  for  hydrogen,  and  state  the  result. 

B.  Dilute,  (a)  Recall  the  litmus  test.  State  it.  Prepare 
some  very  dilute  nitric  acid  by  pouring  a  few  drops  of  the 
ordinary  dilute  acid  into  a  test  tube  half  full  of  water,  dip  a 
glass  rod  into  the  diluted  acid,  and  touch  the  rod  very  cau- 
tiously to  the  tongue.  State  the  result. 

(b)  Recall,  perform,  or  repeat  (if  necessary)  one  or  more 
experiments  illustrating  the  formation  of  salts  of  nitric  acid. 
State  the  results  of  these  experiments. 

(c)  Add  dilute  nitric  acid  to  zinc,  to  copper,  and  to  mag- 
nesium, as  in  A  (/).     State  the  results. 

Required  Exercises. —  i.  What  property  of  nitric  acid  was  shown 
by  A  (6)?  By  (<*)?  By  (e)? 

2.  How  does  the  action  in  A  (b)  and  (c)  compare  with  that  of  hydro- 
chloric acid  under  similar  circumstances? 

3.  Apply  question  2  to  A  (/ ).     (If  in  doubt,  try  the  experiment.) 

Experiment  49  —  Test  for  Nitric  Acid  and  Nitrates 

MATERIALS.  —  Concentrated  nitric  and  sulphuric  acids,  ferrous  sul- 
phate, sodium  nitrate. 

A.  To  a  test  tube  one-fourth  full  of  water  add  a  little  con- 
centrated nitric  acid  and  shake.  Add  an  equal  volume  of  con- 
centrated sulphuric  acid.  Shake  until  the  acids  are  well 
mixed,  then  cool  by  holding  the  test  tube  in  running  water. 
Make  a  cold,  dilute  solution  of  fresh  ferrous  sulphate,  and 


6o 


CHEMISTRY 


pour  this  solution  carefully  down  the  side  of  the  test  tube  upon 
the  nitric  acid  mixture.  Where  the  two  solutions  meet,  a 
brown  or  black  layer  will  appear,  consisting  of  a  compound 
formed  by  the  interaction  of  the  nitric  acid  and  the  ferrous 
sulphate. 

B.  This  test  can  also  be  used  for  a  nitrate.  Proceed  as  above 
with  a  concentrated  solution  of  sodium  nitrate  in  place  of 
nitric  acid.  Record  the  result. 

Experiment  50  —  Preparation  and  Properties  of  Nitric 
Oxide  and  Nitrogen  Dioxide 

MATERIALS.  —  10  grams  of  copper  (borings  or  fine  pieces  of  sheet 
metal),  concentrated  nitric  acid,  pneumatic  trough  filled  as  usual, 
three  bottles,  three  glass  plates,  matches,  piece  of  copper  wire 
(15  centimeters  or  6  inches  long);  and  the  apparatus  shown  in 
Fig.  127. 

Put  the  copper  into  the  bottle,  and  arrange  the  apparatus 
to  collect  the  gas  over  water  (Fig.  127).  Adjust  the  delivery 

tube,  fill  three  bot- 
tles with  water,  and 
invert  them  in  the 
trough.  Dilute  25 
cubic  centimeters  of 
concentrated  nitric 
acid  with  an  equal 
volume  of  water, 
and  introduce  just 
enough  of  this  dilute 
acid  through  the 
dropping  tube  into 
the  bottle  to  cover 
the  copper.  If  the 
action  is  too  vigor- 


Fig.  127.  —  Apparatus  for  Preparing 
Nitric  Oxide. 


ous,  add  water  through  the  dropping  tube;  if  too  weak,  add 
a  little  of  the  dilute  nitric  acid.  Collect  three  bottles  of 
the  gas.  Cover  them  with  glass  plates  and  stand  them  aside 
until  needed. 


NITROGEN— NITROGEN    COMPOUNDS          61 

Proceed  with  the  nitric  oxide  as  follows :  — 

(a)  Observe  its  general  properties  while  covered. 

(b)  Uncover  a  bottle.     Observe  the  result.     Is  the  brown 
gas,  which  is  formed,  identical  in  color  with  the  one  observed 
in  the  generator  at  the  beginning  of  the  experiment? 

(c)  Uncover  a  bottle,  let  the  brown  gas  form,  then  pour  in 
about  25  cubic  centimeters  of  water,  cover  with  the  hand  and 
shake  vigorously,  still  keeping  the  bottle  covered.    Why  does 
the  brown  gas  disappear? 

(d)  With  the  third  bottle,   determine  whether  the  gases 
will  burn  or  support  combustion.     A  convenient  flame  is  a 
burning  match  fastened  to  a  copper  wire.    Plunge  it  quickly 
to  the  bottom  at  first  and  gradually  raise  it  into  the  brown 
gas.     State  the  result. 

NOTE.  —  As  soon  as  (d)  is  performed,  filter  the  blue  liquid  in  the 
generator  bottle,  and  save  the  filtrate  for  Exp.  51. 

Required  Exercises.  —  i.     Summarize  the  properties  of  nitric  oxide. 
Of  nitrogen  dioxide. 

2.  What  is  the  general  chemical  relation  of  the  two  gases  to  each 
other?     To  the  air? 

3.  Why  cannot  nitrogen  dioxide  be  collected  by  displacement  of 
water? 

Experiment  51  —  Properties  of  Nitrates 

MATERIALS.  —  Copper  nitrate  (or  the  solution   from    Exp.  50),  lead 

nitrate. 

Pour  about  50  cubic  centimeters ,  of  the  filtrate  from 
Exp.  50  into  an  evaporating  dish,  stand  the  dish  on  a  gauze- 
covered  ring  and  evaporate  the  solution  (in  the  hood)  to 
about  half  the  original  volume.  Set  the  solution  aside  to 
crystallize,  and  meanwhile  perform  (b).  When  the  crystals 
have  formed,  or  as  soon  after  as  convenient,  remove  them,  and 
dry  them  by  pressing  between  filter  paper.  If  the  filtrate  from 
Exp.  50  was  not  saved,  use  copper  nitrate  from  the  laboratory 
bottle. 

(a)  Put  a  little  of  the  copper  nitrate  in  a  test  tube,  attach 
the  holder,  heat  gently,  and  observe  the  result,  especially  the 


62  CHEMISTRY 

color  of  the  gaseous  product  and  of  the  final  solid  product. 
Test  the  gaseous  product  for  oxygen;  state  the  result.  Devise 
an  experiment  to  determine  the  qualitative  composition  of 
the  solid  product;  submit  the  details  to  the  Teacher  before 
proceeding. 

(b)  Pulverize  a  little  lead  nitrate  and  heat  it  in  a  test  tube 
as  in  (a).  State  the  results. 

SUPPLEMENTARY  EXPERIMENTS 

Experiment  52  —  Preparation  of  Nitrogen  from  Various 
Substances 

MATERIALS.  —  Ammonium  dichromate,  sand,  sodium  nitrate,  potas- 
sium nitrate,  barium  nitrate,  powdered  iron. 

Prepare  nitrogen  by  one  or  more  of  the  following  methods:  — 

A.  Ammonium  Dichromate.     Put  about   2  grams  of  ammonium 
dichromate  in  a  test  tube,  add  5  grams  of  dry  clean  sand,  and  mix  the 
two  substances  thoroughly  by  shaking.     Attach  a  test  tube  holder, 
heat  the  mixture  gently,  and  test  the  escaping  gas  for  nitrogen. 
State  the  result. 

B.  Nitrates  and  Iron.     Mix  thoroughly  about  1.5  grams  of  sodium 
nitrate,  1.5  grams  of  potassium  nitrate,  2  grams  of  barium  nitrate, 
and  10  grams  of  iron;  each  substance  must  be  dry  and  powdered.   Put 
the  mixture  in  a  test  tube,  attach  a  test  tube  holder,  spread  the  mix- 
ture along  the  test  tube,  and  heat  gently.    Test  the  escaping  gas  for 
nitrogen.    State  the  result. 

Experiment  53  —  Preparation  of  Ammonia  Gas  from 
Various  Substances 

(Each  pupil  need  not  perform  all  of  this  experiment.) 

MATERIALS.  — •  Gelatin,  soda-lime,  substances  enumerated  in  A  (b), 
ammonium  sulphate,  sodium  hydroxide  solution,  ammoniacal 
liquor. 

A.  Nitrogenous  Substances,  (a)  Mix  a  little  gelatin  and  soda- 
lime  on  a  piece  of  paper,  slip  the  mixture  into  a  test  tube,  attach  a 
test  tube  holder,  heat,  and  test  the  escaping  gas  with  moist  red  litmus 
paper.  State  the  result. 


NITROGEN— NITROGEN   COMPOUNDS          63 

(b)  Repeat  (a),  using  soda-lime  with  hair,  feather,  leather  scraps, 
pieces  of  horn,  or  hide  powder.  Observe  and  state  the  results. 

B.  Ammonium  Salts,     (a)  Dissolve  a  little  ammonium  chloride 
in  water,  add  a  little  sodium  hydroxide  solution,  warm  gently,  and 
test  (cautiously)  the  liberated  gas  by  its  odor.    What  is  the  gas? 

(b)  'Repeat  (a),  using  ammonium  sulphate  and  sodium  hydroxide 
or  potassium  hydroxide  solution.    State  the  result. 

(c)  Mix  and  grind  together  in  a  mortar  a  little  ammonium  sulphate 
and  calcium  oxide  (lime).    Test  (by  the  odor)  the  gaseous  product, 
and  state  the  result. 

C.  Ammoniacal  Liquor.    Add  powdered  calcium  oxide  (lime)  to 
a  test  tube  half  full  of  ammoniacal  liquor,  warm  gently,  and  test  the 
escaping  gas  for  ammonia.    State  the  result. 

Experiment  54  —  Interaction  of  Nitric  Acid  and  Metals 

MATERIALS.  —  Zinc,  copper,  tin,  iron,  concentrated  nitric  acid. 

Stand  four  test  tubes  in  the  test-tube  rack.  Slip  into  one  a  few 
pieces  of  zinc,  into  another  a  small  piece  of  tin,  into  the  third  a  small 
quantity  of  copper  borings,  and  into  the  fourth  a  small  quantity  of 
clean  iron  filings.  Add  to  each  test  tube  in  succession  enough  con- 
centrated nitric  acid  to  cover  the  metal.  Observe  the  changes,  par- 
ticularly (i)  the  vigor  of  the  action,  (2)  the  properties  of  the  solid 
products,  especially  color  and  solubility,  and  (3)  properties  of  the 
gaseous  products.  Tabulate  these  observations. 

Required  Exercise.  —  Name  the  solid  product  of  the  reaction  in 
each  case.  The  gaseous  product. 

Experiment  55  —  Preparation  and  Properties  of  Sodium 
Nitrite 

MATERIALS.  —  5  grams  of  sodium  nitrate,  5  grams  of  lead,  iron  cru- 
cible, glass  rod. 

Heat  the  mixture  of  lead  and  sodium  nitrate  in  an  iron  crucible, 
which  is  supported  on  the  ring  of  an  iron  stand.  Stir  the  melted  mass 
occasionally  with  a  glass  rod,  and  continue  the  heating  until  most  of 
the  lead  has  disappeared.  Allow  the  mass  to  cool,  cover  it  with  water, 
heat  the  water  to  boiling,  let  the  crucible  cool,  and  then  filter;  add  a 
little  water  to  the  residue  in  the  crucible,  boil,  and  filter  this  portion. 
This  operation  extracts  the  sodium  nitrite.  Add  to  the  combined 


64 


CHEMISTRY 


filtrates  several  drops  of  concentrated  sulphuric  acid.  Observe  the 
result.  How  does  the  result  compare  with  the  action  of  concentrated 
sulphuric  acid  on  sodium  nitrate? 

Required  Exercises. —  i.    What  chemical   change  did  the  sodium 
nitrate  undergo? 

2.  What  is  the  test  for  a  nitrite? 

3.  What  is  the  name  of  the  yellowish  residue? 

Experiment  56  —  Preparation  and  Properties  of  Nitrous 

Oxide 

MATERIALS.  —  10  grams  of  ammonium  nitrate,  pneumatic  trough,  wad 
of  iron  thread,  copper  wire,  three  bottles,  three  glass  plates,  sulphur, 
deflagrating  spoon,  joss  stick.    The  apparatus  is  shown  in  Fig.  128. 
The  parts  A,  C,  D,  E  have  been  used  before;   F,  G,  H  are  exactly 
the  same  as  C,  D,  E  respectively;    B  is  a  large  test  tube. 
Construct  and  arrange  the  apparatus  as  shown  in  Fig.  128.    Put 
10  grams  of  ammonium  nitrate  in  the  flask  A .    The  large  test  tube  B 
remains  empty.    The  end  of  H  rests  on  the  bottom  of  the  pneumatic 
trough,  which  is  filled  as  usual.    Be  sure  the  apparatus  is  gas-tight. 


Fig.  128.  —  Apparatus  for  Preparing  Nitrous  Oxide. 

I.  Preparation.  •  Heat  the  flask  gently  with  a  low  flame,  and  read- 
just the  apparatus  if  it  leaks.  The  ammonium  nitrate  melts  at  first 
and  then  appears  to  boil.  Regulate  the  heat  so  that  the  evolution 


NITROGEN  — NITROGEN    COMPOUNDS          65 

of  the  nitrous  oxide  will  be  slow.  Notice  the  fumes  which  form  in  A , 
and  the  liquid  which  collects  in  B.  Prepare  three  bottles  of  nitrous 
oxide,  covering  each  with  a  glass  plate  as  soon  as  removed  from  the 
trough.  When  the  last  bottle  has  been  collected  and  covered,  remove 
the  end  of  the  delivery  tube  from  the  trough. 

Proceed  at  once  with  II. 

II.   Properties.     Test  the  gas  as  follows:  — 

(a)  Allow  a  bottle  to  remain  uncovered  for  a  few  seconds.    How 
does  nitrous  oxide  differ  from  nitric  oxide? 

(b)  Thrust   a  glowing  joss   stick  into   the   same   bottle   of  gas. 
Observe  the  result.       Is  the  gas   combustible?       Does   it   support 
combustion? 

(c)  (i)  Put  a  piece  of  sulphur  in  a  deflagrating  spoon,  light  it,  and 
lower  the  burning  sulphur  at  once  into  another  bottle  of  gas.    Observe 
the  result.      (2)  Twist  one  end  of  the  copper  wire  around  a  wad  of 
iron  thread.      Heat  the  edge  of  the  wad  an  instant  in  the  flame  and 
then  lower  it  quickly  into  a  bottle  of  the  gas.     Observe  the  result. 
Recall  a  similar  experiment  with  oxygen.    Compare  the  two  results. 

Required  Exercises.  —  i.    Describe  briefly  the  preparation  of  nitrous 
oxide. 

2.  Summarize  the  essential  properties  of  nitrous  oxide. 

3.  What  are  the  fumes  noticed  in  A? 

4.  What  in  all  probability  is  the  other  product  (seen  in  B)  of  the 
chemical  change  in  this  experiment?     Could  it  have  been  an  impurity 
in  the  ammonium  nitrate? 

5.  How  could  you  distinguish  ammonium  nitrate  from  other  nitrates? 

6.  How  could  you  distinguish  nitrous  oxide  from  (a)  the  other  oxides 
of  nitrogen,  (b)  air,  (c)  oxygen,  (d)  hydrogen,  (e)  nitrogen,  (/)  carbon 
dioxide? 

7.  Sketch  the  apparatus,  if  time  permits. 


AIR 


Experiment  57  —  Per  Cent  of  Oxygen  and  Nitrogen 
in  Air 

MATERIALS.  —  Solutions  of  pyrogallic  acid  (10  per  cent)  and  potas- 
sium hydroxide  (50  per  cent),  pneumatic  trough  half  filled  with 
water,  250  and  25  cubic  centimeter  graduated  cylinders.  The 
apparatus  (Fig.  129)  consists  of  a  bottle  holding  about  250  cubic 
centimeters  provided  with  a  tightly  fitting  one-hole  rubber  stopper 
through  which  passes  a  glass  plug.  The  plug,  which  is  made  by 
closing  both  ends  of  a  glass  tube  about  10  centimeters  (4  inches) 
long,  should  fit  tight. 

The  volume  of  the  bottle  is  found  thus:  —  Fill  the  bottle  full 
of  water  from  the  pneumatic  trough.  Push  the  stopper  into  the 
bottle  as  far  as  it  will  go,  insert  the  glass  plug  until  the  inner 
end  is  flush  with  the  inner  surface  of  the  stop- 
per, and  then  draw  a  line  around  the  stopper 
with  a  lead  pencil  to  mark  its  position.  Re- 
move the  stopper.  Pour  water  from  the  bottle 
into  the  250  cubic  centimeter  graduate  until 
the  graduate  is  full  (to  the  250  cc.  mark)  or 
the  bottle  is  empty;  read  the  volume.  If  the 
bottle  holds  more  than  250  cubic  centimeters, 
the  last  portion  of  the  water  in  the  bottle  may 
be  poured  into  the  25  cubic  centimeter  gradu- 

Pig  I20> Ap-  ate-     Record  the  total  volume  of  the  bottle  as 

paratus     for  shown  below. 

finding     Per       Measure  exactly  10  cubic  centimeters  of  pyro- 

Cent  of  Oxy-  gayjc  &c[^  [n  tne  25  cubic  centimeter  graduate, 

gen  and   Ni-          ,  .,     .    ,       , ,       ,     , , ,  A  ,  ,  ,  . 

tro  en  in  Air     a         POUr  °  bottle.       Add    2O    cubic 

centimeters   of   potassium   hydroxide    solution, 

insert  the  rubber  stopper  quickly  to  the  proper  mark,  and  then 

push  the  glass  plug  through  the  stopper  until  the  inner  end  is 

flush  with  the  inner  surface  of  the  stopper.    Shake  the  bottle 


AIR  67 

vigorously  a  few  minutes,  and  then  invert  it  and  watch  the 
surface  of  the  liquid  for  bubbles  of  air,  which  will  enter  if  the 
apparatus  leaks.  If  a  leak  is  detected,  ask  the  Teacher  for 
directions  before  proceeding.  If  the  apparatus  is  tight,  con- 
tinue the  shaking  for  about  half  an  hour.  During  this  operation 
the  oxygen  is  absorbed  by  the  solution. 

Place  the  bottle  on  its  side  beneath  the  water  in  the  pneu- 
matic trough,  inclining  it  slightly  so  that  the  lower  edge  of  the 
bottle  rests  upon  the  bottom  of  the  trough  .and  the  hole  in 
the  stopper  is  beneath  the  surface  of  the  water;  grasp  the  bottle 
firmly  by  the  neck  and  stopper,  and  gradually  pull  out  the 
plug,  taking  care  (i)  not  to  pull  out  the  stopper  and  let  any  of 
the  solution  run  out,  and  (2)  to  keep  the  hole  in  the  stopper 
constantly  below  the  surface.  After  the  water  has  stopped 
running  in,  insert  the  plug,  lift  out  the  bottle,  and  measure 
carefully  the  volume  of  the  final  liquid  in  the  bottle. 

Record  and  calculate  as  follows :  — 

(a)  Volume  of  original  solution  .  . 

(b)  Capacity  of  bottle 

(c)  Volume  of  air  taken  (b  —  a) 

(d)  Final  volume  of  liquid 

(e)  Volume  of  water  which  entered  (d  —  a) 

(f)  Per  cent  of  water  which  entered  (e  -f-  c) 

The  per  cent  of  entering  water  equals  the  per  cent  of  gas 
absorbed,  therefore 

(g)  Per  cent  of  oxygen 

(h)  Per  cent  of  nitrogen  (100  —  g) 

NOTE.  —  This  experiment  disregards  the  argon  and  carbon  dioxide 
in  the  air. 

Experiment  58  —  Water  Vapor  in  the  Air 

(a)  Perform,  recall,  or  repeat  (if  necessary)  Exp.  24  (Deli- 
quescence).   What  does  the  result  show  about  the  air? 

(b)  Place  a  piece  of  lime  upon  a  glass  plate  or  a  block  of 
wood  and  let  it  remain  exposed  to  the  air  an  hour  or  more. 
State  the  result.    Does  this  experiment  verify  the  result  in  (a)? 
If  so,  how? 


68  CHEMISTRY 

(c)  Devise  other  experiments  to  show  that  air  contains  water 
vapor.  Submit  the  details  to  the  Teacher  before  performing 
the  experiment. 

Experiment  59  —  Carbon  Dioxide  in  the  Air 

MATERIALS.  —  Calcium   hydroxide   solution,    barium   hydroxide   solu- 
tion, bottle,  air  blast  apparatus  (for  (6)). 

(a)  Pour  25  cc.  of   clear  calcium  hydroxide  solution  into  a 
bottle  and  let  it  stand  exposed  to  the  air  an  hour  or  more. 
Examine  the  surface  of  the  liquid.    State  the  change  that  has 
occurred.    Explain  the  change. 

(b)  If  an  air  blast  apparatus  is  available,  force  air  through  a 
bottle  half  full  of  clear  barium  hydroxide  solution  until  the 
liquid  is  conspicuously  changed.     Describe  and  explain  the 
change. 

(c)  What  do  (a)  and  (b)  show  about  the  air? 

SUPPLEMENTARY  EXPERIMENT 
Experiment  60  —  Testing  Air 

(a)  Apply  Exps.  58  and  59  to  the  air  in  different  parts  of  the  build- 
ing or  to  the  air  outside  the  building.    Start  the  tests  at  the  same  time 
and  obtain  comparable  results.    State  the  results. 

(b)  Apply  Exp.  58  to  the  air  on  several  days,  especially  days 
when  the  weather  varies  considerably. 

(c)  If  a  hygrometer  is  available,  use  it  in  determining  the  relative 
humidity  of  the  air  out  doors  and  inside  the  laboratory. 

(d)  Apply  Exp.  59  to  the  air  in  the  laboratory,  out  doors,  and  in  a 
recitation  room  which  is  in  use.    Proceed  with  the  testing  as  in  (a) 
(this  experiment). 


EQUIVALENT   WEIGHTS 


Experiment  61  —  Equivalent  of  Zinc  to  Hydrogen 

The  object  of  this  experiment  is  to  find  the  number  of  grams  of  zinc 
chemically  equivalent  to  one  gram  of  hydrogen. 

MATERIALS.  — •  The  apparatus  used  in  Exp.  10  I,  zinc,  thermometer, 

barometer. 

Arrange  the  apparatus  (Fig.  130)  to  collect  a  gas  over  water, 
and  have  it  inspected  by  the  Teacher.  Weigh  a  piece  of  zinc 
on  the  accurate 
balance.  Weigh 
between  .45  and  .5 
gm.,  taking  care  to 
weigh  it  exactly. 
Record  the  weight 
at  once  in  the  note- 
book (as  below). 
Put  the  weighed 
zinc  into  the  gene- 
rator  bottle  A. 
Fill  the  bottle  with 
water  and  insert 
the  stopper  with 
all  its  tubes.  Next 


F/g.  130.  —  Apparatus  for  Finding  the  Equiva- 
lent of  Hydrogen  of  Zinc. 


fill  the  remainder  of  the  apparatus  with  water  by  first  filling 
the  cup  with  water  and  then  admitting  it  repeatedly  until  all 
air  is  forced  out  of  the  bottle  and  tubes;  take  care  never  to 
let  the  water  in  B  fall  below  the  lower  opening  of  the  cup. 
Then  fill  a  collecting  bottle  (250  cc.)  with  water  and  invert 
it  upon  the  support  in  the  wooden  trough ;  put  the  end  of  the 
delivery  tube  under  the  support  and  ask  for  a  final  inspection. 
Heat  about  50  cc.  of  dilute  sulphuric  acid  in  a  test  tube.  Fill 
the  cup  and  introduce  the  hot  acid  in  separate  portions  slowly 


70  CHEMISTRY 

into  the  bottle  A ,  taking  the  same  care  as  before.  Hydrogen 
will  be  slowly  liberated,  and  will  collect  in  the  receiving  bottle. 
Let  the  action  continue  until  all  the  zinc  disappears.  Then 
force  over  into  the  receiving  bottle  all  gas  remaining  in  the 
apparatus  by  admitting  water  carefully  as  before.  Lay  a 
piece  of  dry  filter  paper  upon  the  bottom  of  the  bottle,  grasp 
the  bottle  firmly,  carefully  joggle  it  to  dislodge  any  gas  bub- 
bles which  may  be  underneath  the  support,  slide  the  bottle 
from  the  support,  and  lower  it  into  the  water  until  the  water 
is  about  the  same  level  inside  and  outside  the  bottle;  then 
slip  two  pieces  of  filter  paper  beneath  the  bottle,  cover  the 
mouth  firmly,  remove  the  bottle  from  the  trough  and  stand 
it,  right  side  up,  upon  the  table.  Stand  a  thermometer  in  the 
trough.  Fill  a  250  cc.  graduate  exactly  to  the  mark  with 
water,  remove  the  paper  cover  from  the  bottle,  and  very 
carefully  fill  the  bottle  with  water  from  the  graduate;  read 
and  record  (as  V',  below)  the  exact  volume  of  water  added  — 
which  is,  of  course,  the  volume  of  hydrogen  gas  liberated. 
Read  the  thermometer  while  the  bulb  is  in  the  water,  and 
record  the  reading.  Record  the  barometer  reading.  Find  the 
vapor  pressure  corresponding  to  the  recorded  temperature 
(see  Appendix,  §  4),  and  record  it  as  a  below. 

RECORD 


Weight  of  zinc  taken  (Zn)      .      . 
Observed  volume  of  hydrogen  (V) 

Temperature  (t) 

Pressure  (P)      .      .      . 
Vapor  pressure   (a)       .... 
Corrected  volume  of  hydrogen  (V) 
Equivalent  of  zinc  (E) 


Calculation.  Reduce  the  observed  volume  (V)  of  hydrogen 
to  the  volume  (V)  it  would  have  at  o°  C,  760  mm.,  and  dry 
state  by  the  formula  given  in  Part  I,  §  40,  viz.  - 

V'  (P  -  a) 


V  = 


760  (i  +  (.00366  Xt)) 


EQUIVALENT    WEIGHTS  71 

Since  1000  cc.  of  dry  hydrogen  weigh  .0898  gm.,  the  weight 
of  the  corrected  volume  (V)  is  found  by  1000:  V::  .0898:  X. 
And  the  weight  of  zinc  equivalent  (E)  to  one  gram  of  hydrogen 
is  found  by  X:  Zn::  i:  E. 

Submit  the  result  to  the  Teacher  for  criticism  (before  taking 
the  apparatus  apart,  if  convenient). 

Experiment  62  —  Equivalent  of  Magnesium  to  Hydrogen 

MATERIALS.  —  A  100  cc.  tube  and  wooden  trough,  magnesium  ribbon, 
thermometer,  barometer. 

Weigh  accurately  between  .065  and  .075  gm.  of  magnesium 
ribbon,  preferably  in  a  single  piece.  Have  the  wooden  trough 
half  full  of  water.  Pour  8  cc.  of  concentrated  hydrochloric 
acid  into  the  100  cc.  tube  and  fill  the  tube  completely  with  cold 
water.  Put  the  magnesium  into  the  tube,  cover  the  end  of 
the  tube  with  the  thumb  or  finger,  invert  the  tube,  stand  it  in 
the  trough,  but  keep  the  end  loosely  closed  to  prevent  the 
magnesium  from  slipping  out.  As  the  acid  reaches  the  mag- 
nesium, action  begins  vigorously.  Hydrogen  rises  in  the  tube 
and  usually  carries  the  magnesium  with  it.  Watch  the  opera- 
tion, and  agitate  the  tube  to  prevent  the  magnesium  from 
sticking  to  the  inside.  The  action  is  very  rapid  and  must  be 
watched  constantly,  being  over  in  about  two  minutes.  If  a 
piece  of  magnesium  should  stick  to  the  inside  of  the  tube, 
close  the  end  of  the  tube  tightly,  lift  it  from  the  water,  incline 
it  enough  to  loosen  the  magnesium,  and  then  quickly  straighten 
the  tube  and  put  the  end  beneath  the  water. 

When  all  the  magnesium  has  disappeared,  close  the  end  of 
the  tube,  remove  the  tube  to  a  tall  jar  of  water,  and  let  it  stand 
five  minutes;  then,  without  touching  the  tube  with  the  bare 
hands,  adjust  the  height  so  that  the  water  levels  are  the  same 
inside  and  outside  of  the  tube,  and  read  the  volume.  Read 
the  barometer  and  the  thermometer  (keeping  the  bulb  in  the 
water).  Calculate  the  equivalent  of  magnesium  as  in  Exp.  61. 


72  CHEMISTRY 

SUPPLEMENTARY  EXPERIMENTS 
Experiment  63  —  Equivalent  of  Iron  to  Copper 

The  object  of  this  experiment  is  to  find  the  weight  of  copper  precipi- 
tated by  a  known  weight  of  iron. 

MATERIALS.  —  Beaker,  glass  rod,  iron  powder,  copper  sulphate  solu- 
tion of  known  strength,  alcohol. 

Prepare  or  obtain  about  50  cc.  of  a  copper  sulphate  solution  which 
contains  .1  gm.  of  copper  to  i  cc.  Weigh  a  clean  dry  beaker.  Weigh 
it  in  accurately  about  2  gm.  of  iron  powder.  Add  slowly  about  25  cc. 
of  the  copper  sulphate  solution.  The  iron  precipitates  the  copper  as  a 
fine  powder.  Stir  occasionally  with  the  glass  rod.  About  one  hour  is 
needed  for  complete  precipitation.  When  it  is  judged  that  all  the  iron 
has  been  used  up,  let  the  copper  settle,  and  pour  off  the  liquid  down 
the  rod  into  a  dish,  taking  care  not  to  lose  any  copper.  Add  water 
to  the  beaker,  stir,  let  settle,  and  decant  as  before.  If  the  wash  water 
contains  particles  of  copper,  let  it  settle,  pour  off  the  water  and  add 
the  copper  to  the  beaker.  Wash  until  the  washings  give  no  test  for  a 
sulphate  (i.e.  no  white  precipitate  of  barium  sulphate  upon  addition 
of  barium  chloride).  Finally  add  a  little  alcohol  and  heat  to  dryness 
very  cautiously  on  a  piece  of  asbestos.  When  dry  and  cool,  weigh 
quickly  before  the  copper  oxidizes. 

Calculation.  Assume  31.8  as  the  equivalent  of  copper  and  calcu- 
late the  equivalent  of  iron. 

Experiment  64  —  Equivalent  of  Aluminium  to  Hydrogen 
MATERIALS.  —  As  in  Exp.  61. 

Proceed  as  in  Exp.  61  (Equivalent  of  Zinc),  but  (a)  weigh  out 
about .  1 7  gm.  of  aluminium  (taking  care  to  weigh  exactly  the  amount 
used)  and  (b)  use  hot  concentrated  hydrochloric  acid  instead  of 
dilute  sulphuric  acid.  Record  and  calculate  as  in  Exp.  61. 

Experiment  65  —  Equivalent  of  Calcium  to  Hydrogen 
MATERIALS.  —  As  in  Exp.  62. 

Proceed  as  in  Exp.  62,  but  use  about  .115  gm.  of  calcium.  Record 
and  calculate  as  in  Exp.  62. 


SOLUTION  — ACIDS,   BASES,   SALTS 

Experiment  66  —  Chemical  Behavior  of  Electrolytes 
in  Solution 

MATERIALS.  —  Solutions  of  silver  nitrate,  hydrochloric  acid,  ammonium 
chloride,  barium  chloride,  calcium  chloride,  magnesium  chloride, 
sodium  chloride,  potassium  chloride,  potassium  chlorate,  potassium 
perchlorate,  chloroform. 

(a)  Test  separately  dilute  solutions  of  each  of  the  following 
substances  for  ionic  chlorine  (i.e.  for  chloride  ions)  by  adding 
a  few  drops  of  silver  nitrate  solution,  and  state  the  result  in 
each  case:    Hydrochloric  acid,  ammonium  chloride,  barium 
chloride,    calcium     chloride,     magnesium    chloride,    sodium 
chloride,  potassium  chloride. 

(b)  Test  a  solution  of  potassium  chlorate  for  chloride  ions. 
State  the  result. 

(c)  Repeat  (b),  using  a  solution  of  potassium  perchlorate 
instead  of  potassium  chlorate.    State  the  result. 

(d)  Shake  a  little  chloroform  with  water,  and  test  as  in  (b). 
State  the  result. 

Required  Exercises. —  i.    What  ions  are  in  a  solution  of  all  chlorides? 

2.  What  ions  are  in  a  solution  of  potassium  chlorate?     Of  silver 
nitrate? 

3.  Explain  the  general  result  in  (a)  and  the  results  in  (b),  (c),  and  (d) 
in  terms  of  the  theory  of  electrolytic  dissociation. 

Experiment  67  —  Chemical  Behavior  of  Electrolytes 
in  Solution 

MATERIALS.  —  Solutions  of  barium  chloride,  sulphuric  acid,  copper 
sulphate,  sodium  sulphate,  aluminium  sulphate,  magnesium  sul- 
phate, zinc  sulphate. 

Test  dilute  solutions  of  the  following  for  sulphate  ions  by 
adding  to  each  separately  a  few  drops  of  barium  chloride 
solution,  and  state  the  result  in  each  case:  Sulphuric  acid, 


74  CHEMISTRY 

copper  sulphate,  sodium  sulphate,  aluminium  sulphate,  mag- 
nesium sulphate,  zinc  sulphate. 

Required  Exercises.  —  i.   What  ions  are  in  solutions  of  all  sulphates? 
2.    Explain  the  general  result  obtained  above  in  terms  of  the  theory 
of  electrolytic  dissociation. 

Experiment  68  —  General  Properties  of  Acids,  Bases, 

and  Salts 
MATERIALS.  —  As  in  Exps.  37,  38,  39,  40. 

Recall  and  state  the  general  properties  of  acids,  bases,  and 
salts,  or  repeat  (if  necessary)  Exps.  37,  38,  39,  40. 

Experiment  69  —  The  Litmus  Reaction  of  Different  Salts 

MATERIALS.  —  Acid  sodium  phosphate,  acid  sodium  sulphate,  potas- 
sium (or  sodium)  nitrite,  sodium  acetate,  potassium  iodide,  sodium 
carbonate,  potassium  carbonate,  potassium  dichromate,  sodium 
sulphate,  copper  sulphate,  ferric  chloride. 

Prepare,  or  obtain,  dilute  solutions  of  the  salts  .mentioned 
above,  and  test  them  with  both  kinds  of  litmus  paper. 

Required  Exercises. —  i.  Classify  these  salts  under  the  headings 
Normal,  Acid,  and  Basic. 

2.  Name  the  salts  used  above  that  .undergo  hydrolysis.  Interpret 
hydrolysis  by  the  theory  of  electrolytic  dissociation. 

Experiment  70  —  Electrolysis  of  Copper  Sulphate 

MATERIALS.  —  Dilute  solution  of  copper  sulphate,  small  battery  jar 
(or  beaker),  two  electrodes  (pieces  of  electric  light  carbon)  and 
connection  wires,  battery  of  four  or  more  cells  (or  other  source  of 
electric  current). 

Fill  the  battery  jar  about  two-thirds  full  of  dilute  copper  sul- 
phate solution.  Wind  one  end  of  a  piece  of  the  wire  around 
one  end  of  each  electrode  and  hang  the  electrodes  in  the  solu- 
tion by  bending  the  wire  over  the  edge  of  the  jar  (or  by  the 
device  shown  in  Fig.  32,  Part  I).  Before  turning  on  the  cur- 
rent (or  making  the  final  connection),  examine  each  electrode 
and  note  the  absence  of  any  deposit.  (After  the  apparatus  is 


SOLUTION  —  ACIDS,  BASES,  SALTS 


75 


set  up,  the  Teacher  should  mark  the  anode  and  cathode.) 
Turn  on  the  current  and  observe  what  takes  place  at  the  anode. 
When  the  current  has  run 
about  ten  minutes,  shut  it 
off,  and  examine  each  elec- 
trode. Compare  with  their 
appearance  before  the  elec- 
trolysis took  place.  Upon 
which  electrode  is  there  a 
deposit?  What  is  the  de- 
posit? 

Sketch    the    apparatus,     F[^   131-- Apparatus  for  Showing 
.,  ,  the  Behavior  of  Solutions  toward 

and     describe     the     elec-        an  Electric  Current. 

trolysis  of  copper  sulphate 

in  terms  of  the  theory  of  electrolytic   dissociation,  using  the 

sketch  in  your  interpretation. 


Experiment  71  —  Electrolysis  of  Sodium  Sulphate 

MATERIALS.  —  Sodium  sulphate  solution>  litmus  solution  (preferably 
neutral),  U-tube  clamped  to  an  iron  stand,  narrow  aluminium 
electrodes  (to  fit  the  U-tube)  and  connection  wires,  battery  of  four 
or  more  cells  (or  other  source  of  electric  current). 

Fill  the  U-tube  two-thirds  full  of  sodium  sulphate  solution, 
and  add  enough  litmus  solution  to  produce  a  faint  color  after 
shaking.  Attach  the  U-tube  to  the  iron  stand,  insert  the 
electrodes,  and  note  the  color  of  the  solution  in  each  arm  of 
the  U-tube.  (After  the  apparatus  is  set  up,  the  Teacher  should 
mark  the  anode  and  cathode.)  Turn  on  the  current  and  let  it 
run  until  there  is  a  change  in  color  in  each  arm  of  the  U-tube. 
Note  this  color  and  note  also  whether  gas  is  liberated  in  each 
arm  of  the  U-tube. 

Sketch  the  apparatus  (except  the  battery)  and  interpret 
the  electrolysis  of  sodium  sulphate  by  the  theory  of  electrolytic 
dissociation,  using  the  sketch  in  your  interpretation. 


76  CHEMISTRY 

SUPPLEMENTARY  EXPERIMENTS 
Experiment  72  —  Electrolytes  and  Non-Electrolytes 

MATERIALS.  —  See  Part  I,  §  156.    The  apparatus  is  shown  in  Fig.  131. 
Proceed  with  different  solutions  as  described  in   Part  I,   §  156. 
Tabulate  the  results. 

Experiment  73  —  Chemical  Behavior  of  Electrolytes  in 
Solution 

MATERIALS.  —  Solutions  of  silver  nitrate,  silver  sulphate,  potassium 
chloride,  barium  chloride,  barium  nitrate,  potassium  sulphate. 

A.  (a)  Add  a  few  drops  of  potassium  chloride  solution  to  a  little 
silver  nitrate  solution.    Shake  well  and  observe  the  result. 

(b)  Repeat  (a),  using  potassium  chloride  and  silver  sulphate  solu- 
tions. Compare  the  results.  Are  the  precipitates  identical? 

B.  (a)  Proceed  as  in  A  (a)  with  potassium  sulphate  and  barium 
chloride  solutions.    Observe  the  result. 

(6)  Proceed  as  in  B  (a)  with  potassium  sulphate  and  barium  nitrate 
solutions.  Compare  the  result  with  that  in  B  (a).  Are  the  precipi- 
tates identical? 

Required  Exercises.  —  i.  What  have  solutions  of  silver  nitrate  and 
silver  sulphate  in  common?  Solutions  of  barium  chloride  and  barium 
nitrate? 

2.    Name  the  ionic  substances  in  the  solutions  in  A.     In  B. 

Experiment  74.  —  Electrolysis  of  Potassium  Iodide 

MATERIALS.  —  Starch,    potassium    iodide,    mortar    and    pestle,    filter 
paper,  sheet  of  metal  (tin  or  iron),  battery  of  two  or  more  cells. 

Grind  together  in  a  mortar  a  lump  of  starch  and  a  crystal  of  potas- 
sium iodide.  Add  enough  water  to  make  a  thin  liquid.  Dip  a  strip 
of  filter  paper  into  the  mixture,  and  spread  the  wet  paper  upon  the 
sheet  of  metal.  Press  the  end  of  the  wire  attached  to  the  zinc  (of  the 
battery)  upon  the  metal,  and  draw  the  other  wire  across  the  sheet  of 
paper.  The  marks  are  caused  by  iodine  which  is  liberated  from  the 
potassium  iodide  and  colors  the  starch. 

Required  Exercises.  —  i.    Describe  briefly  this  experiment. 

2.  Iodine  is  a  non-metal.  Are  iodine  ions  anions  or  cations?  At 
what  electrode  is  iodine  liberated? 


SOLUTION  —  ACIDS,  BASES,   SALTS  77 

Experiment  75  —  Electrolysis  of  Water 

Recall  the  experiment  showing  the  electrolysis  of  water  (see 
Exp.  30). 

Required  Exercises.  —  i.  State  briefly  the  explanation  of  the  elec- 
trolysis of  water  in  terms  of  the  theory  of  electrolytic  dissociation. 

2.  Are  hydrogen  ions  anions  or  cations?    To  what  electrode  do  hydro- 
gen ions  migrate? 

3.  Is  oxygen  a  primary  or  a  secondary  product  of  the  electrolysis 
of  water? 

4.  If  oxygen  ions  were  formed  in  the  solution,  (a)  would  they  be 
anions  or  cations,  and  (6)  to  what  electrode  would  they  migrate? 

Experiment  76  —  Colored  Ions 

MATERIALS.  —  Copper  sulphate,  copper  nitrate,  copper  bromide,  nickel 
chloride,  nickel  sulphate,  cobalt  chloride,  cobalt  nitrate,  potas- 
sium dichromate,  ammonium  dichromate,  sodium  dichromate, 
potassium  chromate,  potassium  permanganate. 

A.  Copper  Ions.     Prepare  a  dilute  solution  of  each  of  the  copper 
compounds  mentioned  above  by  dissolving  a  little  of  the  solid  in  a 
test  tube  half  full  of  water.    Compare  the  colors. 

B.  Nickel  Ions.     Prepare,  .or  obtain  from  the  Teacher,  a  dilute 
solution  of  each  of  the  nickel  compounds,  and  compare  the  colors. 

C.  Cobalt  Ions.     Proceed  as  in  B  with  the  cobalt  compounds. 

D.  Miscellaneous.     Determine  the  color  of  dichromate  ions.     Of 
chromate  ions.    Of  permanganate  ions. 

Required  Exercise.  —  Name  several  kinds  of  colorless  ions. 

Experiment  77  —  Migration  of  Ions 

MATERIALS.  —  Battery  (or  other  source  of  an  electric  current),  U-tube, 
two  strips  of  aluminium  to  fit  the  U-tube,  potassium  dichromate 
solution,  copper  sulphate  solution. 

A.  Potassium  Dichromate.  Fill  a  U-tube  two  thirds  full  of  dilute 
potassium  dichromate  solution  and  clamp  it  in  an  upright  position 
to  an  iron  stand.  Insert  the  electrodes  and  allow  the  current  to  flow 
about  ten  minutes.  Observe  the  color  of  the  solution  when  the  cur- 
rent starts,  the  gradual  change  in  color  in  each  arm  of  the  U-tube 
as  the  current  continues,  and  the  difference  in  color  in  the  two  arms 
when  the  current  stops. 


CHEMISTRY 


B.   Copper  Sulphate.     Proceed  as  in  A  with  dilute  copper  sulphate 
solution. 

Required  Exercises. —  i.    Describe  the  experiment  and  sketch  the 
apparatus. 

2.  Give  the  name  and  formula  of  all  the  ions  and  state  the  electrodes 
to  which  each  kind  of  ion  migrates. 

3.  Give  the  name  and  formula  of  the  colored  ions. 

Experiment  78  —  Neutralization  by  Titration 

The  object  of  this  experiment  is  to  find  the  number  of  grams  of  the 
compound  HC1  in  i  cubic    centimeter  of    a  solution  of  hydrochloric 
acid  (i.e.  HC1  dissolved  in  water)  by  neutralizing  the  acid  with  a  solu- 
tion of  sodium  hydroxide  of  known  concentration. 
MATERIALS.  —  Two    burettes     (Fig.    132),     two   beakers,   a    glass    rod, 
phenol-phthalein  solution,  and  solutions  of  hydrochloric  acid  and 
sodium  hydroxide  (the  latter  of  known  concentration  and  obtained 
from  the  Teacher). 

Fill  each  burette  (or  start  with  each  full)  —  one  with  the  acid 
solution  and  one  with  the  base  solution,  as 
marked.  Place  the  waste  beaker  under  each 
burette  in  turn  and  allow  the  solution  to  run 
out  slowly  until  the  bottom  of  the  meniscus 
rests  on  the  O  line  when  the  eye  is  on  a  level 
with  the  same  line.  (See  Fig.  132.)  Set  the 
waste  beaker  aside.  Put  a  clean  beaker 
under  the  base  burette  and  let  exactly  15 
cc.  run  into  the  beaker;  record  as  in  I  be- 
low. Remove  the  beaker,  add  2  or  3  drops 
of  phenol-phthalein  solution,  put  the  beaker 
under  the  acid  burette  and  let  the  acid  solu- 
tion run  in  slowly,  stirring  constantly  with 
the  clean  rod  until  the  red  color  just  dis- 
appears and  the  solution  becomes  colorless. 
-Burettes.  Read  ^  exact  volume  of  add  solution 

(Enlarged Section  (on       ,,    ,        ,  ,        .     T      _  .•    . 

\    01  .1  added  and  record  as  in  I.    Pour  the  solution 

page    9)    Shows    the  . 

Curved     Surface    of  °  beaker,  wash  the  beaker,  and  pro- 

the     Solution.     Cor-  ceed>    as    before,   with    a    second    15  cc.  of 

rect  Reading  of  the  NaOH  solution.    Record  as  in  II.    Wash  the 

Volume  of  the  Solu-  beaker  and  proceed   with   a  third  15  cc.  of 

tion  is  along  Line  I.)  NaOH  solution.    Record  as  in  III. 


SOLUTION  —  ACIDS,  BASES,  SALTS      79 

RECORD 

I.     NaOH  sol.  o    -15  =  15 

HC1    "    o  '-  i  cc.  NaOH  sol.  =     cc.  HC1  sol. 

II.     NaOH  sol.  15  -  30  =  15 

HC1     "  i  cc.  NaOH  sol.  =     cc.  HC1  sol. 

III.     NaOH  sol.  30  -  45  =  i5 

HC1     "         -  i  cc.  NaOH  sol.  =     cc.  HC1  sol. 

Calculation:  —  (a)  Write  the  equation,  including  the  weights  of 
the  NaOH  and  HC1. 

(b)  Find  (from  I,  II,  III)  the  average  number  of  cc.  of  HC1  solution 
equal  to  i  cc.  of  NaOH  solution. 

(c)  i  cc.  of  NaOH  solution  contains  ?  gm.  (ask  Teacher)  NaOH. 

(d)  From  the  value  in  (c)  and  the  relative  values  of  the  solutions 
together  with  the  values  in  the  equation,  calculate  the  number  of 
gm.  of  the  compound  HC1  in  i  cc.  of  the  acid  solution.    Submit  the 
result  to  the  Teacher  for  criticism. 

Experiment  79  —  Preparation  of  Salts 

(Each  pupil  need  not  perform  all  of  this  experiment.) 

MATERIALS.  —  Calcium,  calcium  oxide,  calcium  carbonate,  calcium 
chloride,  silver  nitrate  solution,  evaporating  dish,  gauze-covered 
ring. 

A.  Acid  and  a  Metal.  Put  a  small  piece  of  clean  calcium  in  an 
evaporating  dish,  add  a  little  dilute  hydrochloric  acid,  stand  the  dish 
on  a  gauze-covered  ring,  and  heat  the  dish  gently  until  the  calcium 
disappears,  adding  more  acid  if  necessary.  Then  evaporate  the  solu- 
tion to  dryness  in  the  hood,  taking  care  to  heat  gently  toward  the  end 
of  the  evaporation  to  prevent  spattering.  Add  just  enough  water  to 
moisten  the  residue,  and  evaporate  again  to  dryness.  Heat  the  residue 
until  no  more  fumes  of  hydrochloric  acid  are  evolved.  Let  the  dish 
cool,  and  loosen  the  contents  with  a  glass  rod.  Test  portions  of  the 
residue  for  (a)  calcium  and  (b)  a  chloride  as  follows:  (a)  Touch  a 
clean,  moist  test  wire  to  a  small  piece  of  the  residue,  and  hold  it  in 
the  outer  and  lower  edge  of  the  Bunsen  flame.  The  yellow-red  color 
imparted  to  the  flame  is  caused  by  the  calcium,  and  is  one  test  for 
this  element,  (b)  Dissolve  a  little  of  the  residue  in  a  test  tube  half 
full  of  water,  and  apply  the  usual  test  for  a  chloride  to  this  solution. 
State  the  result. 


8o  CHEMISTRY 

Required  Exercises. —  i.  What  is  the  name  and  formula  of  the 
residue  formed  in  this  experiment?  Write  the  equation  for  the  reaction. 

2.  Suggest  an  experiment  to  verify  the  answer  to  i. 

3.  Suggest  experiments  to  prove  that  the  compound  is  neither  an 
acid  nor  a  base. 

4.  Cite  two  or  more  experiments,  already  performed,  which  illus- 
trate this  method  of  salt  formation. 

B.  Acid  and  an  Oxide.    Proceed  as  in  A,  using  hydrochloric  acid 
and  a  small  piece  of  calcium  oxide.    Before  evaporating  to  dryness 
filter  the  solution,  if  it  is  not  clear.    Test  the  final  residue  as  in  A 
and  .state  the  result. 

Required  Exercises.  —  As  in  A  (except  4). 

C.  Acid  and  a  Salt,    (a)  Proceed  as  in  A,  using  hydrochloric  acid 
and  several  small  pieces  of  calcium  carbonate.    Test  the  final  residue 
(as  in  A)   obtained  by  evaporating  the  clear  solution.     State  the 
result. 

(6)  Put  a  little  sodium  chloride  in  an  evaporating  dish,  add  25  cc. 
of  dilute  sulphuric  acid,  stand  the  dish  on  a  gauze-covered  ring,  and 
evaporate  the  solution  to  dryness  in  the  hood.  Toward  the  end  of 
the  evaporation,  it  may  be  necessary  to  remove  the  burner,  turn  down 
the  flame,  and  heat  very  gently  by  moving  the  burner  slowly 
back  and  forth  beneath  the  dish.  As  soon  as  the  danger  of  spattering 
is  over,  heat  strongly  as  long  as  white  choking  fumes  are  evolved;  this 
operation  removes  the  last  portions  of  sulphuric  acid  and  completes 
the  chemical  change.  Let  the  dish  stand  on  the  gauze  until  cool 
enough  to  handle.  Then  remove  it,  and  loosen  the  solid  with  a  glass 
rod  or  a  knife.  Test  portions  of  the  residue  for  sodium  and  a  sul- 
phate as  follows:  Proceed  with  the  flame  test  for  sodium  as  in  A  (a), 
and  observe  and  state  the  result.  Apply  the  usual  test  for  a  sulphate 
to  a  solution  of  the  residue,  and  state  the  result. 

Required  Exercises.  —  i.    For  (a),  as  in  A. 

2.    For  (&),  as  in  A  i  and  3. 

D.  Two  Salts.  Add  a  little  silver  nitrate  solution  to  a  little  cal- 
cium chloride  solution,  and  describe  the  result. 

Required  Exercises.  —  As  in  A  (except  4). 

IE.  Acid  and  Base.  Recall  two  or  more  experiments  in  which  a 
salt  was  formed  by  the  interaction  of  an  acid  and  a  base,  and  name 
all  the  compounds  involved  in  the  reactions.  Write  the  equation  for 
reaction  in  (a)  the  ordinary  form  and  (b)  the  ionic  form. 


CARBON 

Experiment  80  —  Distribution  of  Carbon 

MATERIALS.  — •  Sand    (or   clay)    crucible,   sand,   wood,   cotton,   starch, 
sugar,  glass  tube  (or  rod),  candle,  block  of  wood. 

(a)  Cover  the  bottom  of  the  crucible  with  a  thin  layer  of 
sand.    Put  on  the  sand  a  small  piece  of  wood,  a  small,  compact 
wad  of  cotton,  and  a  lump  of  starch.    Fill  the  crucible  loosely 
with  dry  sand,  and  slip  it  into  the  ring  of  an  iron  stand.    Heat 
with  a  flame  which  extends  well  above  the  bottom  of  the 
crucible  until  the  smoking  ceases  (approximately  20  minutes). 
After  the  crucible  has  cooled  sufficiently  to  handle,  pour  the 
contents  out  upon  a  block  of  wood  or  an  iron  pan.    Examine 
the  contents.     What  is  the  residue?    What  is  hereby  shown 
about  the  distribution  of  carbon? 

While  the  crucible  is  heating,  proceed  as  follows: 

(b)  Heat  a  little  sugar  in  a  test  tube  until  the  vapors  cease 
to  appear.    What  is  the  most  obvious  product? 

(c)  Close  the  holes  at  the  bottom  of  a  lighted  Bunsen  burner, 
and  hold  a  glass  tube  in  the  upper  part  of  the  flame  long  enough 
for  a  thin  deposit  to  form.     Examine  it,  name  it,  and  state 
its  source. 

(d)  Hold  a  glass  tube  in  the  flame  of  a  candle  which  stands 
on  a  block  of  wood,  and  compare  the  result  with  that  in  (c). 

Experiment  81  —  Properties  of  Coal 

MATERIALS.  —  Anthracite    and    bituminous    coal,    lignite,    crucible, 
large  graduated  cylinder. 

(#)  Examine  specimens  of  anthracite  and  bituminous  coal, 
and  lignite,  and  state  the  characteristic  properties,  e.g.  rela- 
tive hardness,  color,  luster. 

(b)  Pulverize  a  little  of  each  variety,  heat  gently  at  first  in 
a  crucible  or  a  test  tube,  and  observe  the  result,  especially  the 


82  CHEMISTRY 

liberation  of  carbonaceous  volatile  matter  and  moisture;  then 
heat  strongly  and  observe  the  immediate  result;  continue  the 
heating  until  little  or  no  black  residue  remains.  Summarize  the 
results. 

(c)  Determine  the  specific  gravity  of  coal  by  the  method 
given  in  Exp.  88  (b).     State  the  result. 

(d)  If  fossils  from  a  coal  bed  are  available,  examine  and 
describe  them. 

Experiment  82  —  Properties  of  Charcoal 

MATERIALS.  —  Wood  charcoal  (lump  and  powder),  animal  charcoal, 
copper  wire,  crucible,  vinegar,  hydrogen  sulphide  solution,  test 
tube  fitted  with  a  cork. 

(a)  Examine  a  typical  specimen  of  wood  charcoal  and  state 
its  characteristic  properties.     Do  the  same  with  animal  char- 
coal.    Put  a  little  animal  charcoal  in  a  crucible  and  heat  it 
strongly.    Meanwhile  proceed  with  the  wood  charcoal.    Wind 
the  end  of  a  nichrome  test  wire  around  a  small  piece  of  char- 
coal, hold  it  in  the  flame,  and  observe  the  result,  especially  the 
ease  or  difficulty  of  ignition,  presence  or  absence  of  flame  and 
of  smoke,  formation  of  ash.     Compare  the  results  with  those 
obtained  in  Exp.  81  (b).     When  the  animal  charcoal  has  been 
heated  thirty  or  more  minutes,  examine  the  residue.     What 
is  it? 

(b)  Fill  a  test  tube  one-fourth  full  of  powdered  animal  char- 
coal as  follows:  Fold  a  narrow  strip  of  smooth  paper  so  that  it 
will  slip  easily  into  the  test  tube;  place  the  powder  at  one  end 
of  the  troughlike  holder,  slowly  push  the  paper  into  the  test 
tube,  holding  both  tube  and  paper  in  a  horizontal  position; 
now  hold  the  tube  upright,  and  the  powder  will  slip  from  the 
paper.    Add  10  cubic  centimeters  of  hydrogen  sulphide  solu- 
tion, and  cork  securely.    If  the  tube  leaks,  make  the  opening 
gas-tight   with  vaseline.     Shake   thoroughly.     After   ten   or 
fifteen  minutes,  remove  the  stopper  and  smell  of  the  contents. 
Is  the  odor  much  less  offensive?     Repeat,  unless  a  definite 
result  is  obtained.     What  property  of  animal  charcoal  is  illus- 
trated by  this  experiment? 


CARBON  83 

(c)  Fill  a  test  tube  one  fourth  full  of  powdered  animal  char- 
coal as  in  (b),  add  10  cubic  centimeters  of  vinegar,  shake  thor- 
oughly for  a  minute,  and  then  warm  gently.  Filter  into  a  clean 
test  tube.  Compare  the  color  of  the  filtrate  with  that  of  the 
original  vinegar.  Describe  briefly.  What  property  of  animal 
charcoal  is  illustrated  by  this  experiment? 

Experiment  83  —  Preparation  and  Properties  of  Carbon 
Dioxide 

MATERIALS.  —  Calcium  carbonate,  dilute  hydrochloric  acid,  joss  stick, 
candle  fastened  to  a  wire,  calcium  hydroxide  solution,  four  bottles. 
The  apparatus  is  shown  in  Fig.  133. 

I.  Preparation.  Put  six  or  more  lumps  of  calcium  car- 
bonate into  the  bottle,  and  arrange  the  apparatus  to  collect 


Fig.  133.  —  Apparatus  for  Preparing  Carbon  Dioxide. 

the  gas  over  water,  as  previously  directed.  Introduce  enough 
dilute  hydrochloric  acid  through  the  dropping  tube  to  cover 
the  calcium  carbonate.  Collect  four  bottles,  cover  with  glass 
plates  or  filter  paper,  and  stand  aside  till  needed.  Proceed  at 
once  with  II. 

II.     Properties,     (a)  Plunge   a   blazing   joss   stick   several 
times  into  a  bottle.    State  the  result. 


84  CHEMISTRY 

(b)  Lower  a  lighted  candle  into  a  bottle  of  air,  and  quickly 
invert  a  bottle  of  carbon  dioxide  over  it,  holding  the  bottles 
mouth  to  mouth.    State  the  final  result. 

(c)  Pour  a  little  calcium  hydroxide  solution  into  a  bottle 
of  carbon  dioxide,  cover  it  with  the  hand,  and  shake  it  vigor- 
ously.   Describe  and  explain  the  result. 

(d)  Fill  a  bottle  of  carbon  dioxide  one-third  full  of  water, 
cover  it  tightly  with  the  hand,  and  shake  it  vigorously.    Invert 
the  bottle,  still  covered,  in  the  pneumatic  trough.     Observe 
and  state  the  result. 

NOTE. — As  soon  as  (d)  is  performed  wash  the  acid  from  the  marble 
and  save  the  solid  for  other  experiments. 

Required  Exercises.  —  i.  Describe  briefly  the  preparation  of  carbon 
dioxide. 

2.  What  do  (a)  and  (&)  show  about  the  relation  of  carbon  dioxide 
to  combustion? 

3.  What  does  (b)  show  about  the  relative  weight  of  carbon  dioxide 
and  air? 

4.  What  does  (d)  show  about  the  solubility  of  carbon  dioxide? 

5.  What  is  the  test  for  carbon  dioxide? 

Experiment  84  —  Carbon  Dioxide  and  Respiration 

Exhale  the  breath  through  a  glass  tube  into  a  test  tube  half 
full  of  calcium  hydroxide  solution:  Describe  and  explain  the 
result. 

Experiment  85  —  Preparation  and  Properties  of  Acid 
Calcium  Carbonate 

MATERIALS.  —  Calcium   hydroxide   solution   and   the   carbon    dioxide 
generator  used  in  Exp.  83. 

Pass  carbon  dioxide  into  a  test  tube  half  full  of  calcium 
hydroxide  solution  until  the  precipitate  at  first  formed  dis- 
appears. Filter,  if  the  liquid  is  not  perfectly  clear.  Heat  the 
test  tube  gently  and  observe  carefully  all  the  changes.  State 
the  results  of  heating  the  clear  solution. 

Required  Exercises.  —  i.    What  is  the  name  of  the  first  precipitate? 

2.  Into  what  soluble  compound  was  this  precipitate  formed  by  inter- 
action with  carbon  dioxide? 


CARBON  85 

3.  Into  what  was  the  soluble  compound  formed  by  heating? 

4.  Write  equations  for  the  two  essential  reactions  in  Exp.  85. 

Experiment  86  —  Testing  for  Carbonates 

MATERIALS.  —  Barium  hydroxide  solution,  glass  tube,  baking  soda, 
washing  soda,  baking  powder,  native  chalk,  tooth  powder,  white 
lead,  whiting,  old  mortar  (or  plaster). 

Test  the  substances  enumerated  above  for  the  presence  of 
a  carbonate  as  follows:  Put  a  little  of  the  solid  in  a  test  tube, 
add  a  little  water  and  dilute  hydrochloric  acid,  and  shake; 
then  hold  the  glass  tube,  which  has  been  dipped  into  barium 
hydroxide  solution,  inside  the  test  tube  for  a  minute  or  two 
about  3  centimeters  above  the  mixture.  If  the  action  is  not 
marked,  gently  warm  the  test  tube.  State  the  result  in  each 
case. 

Experiment  87.  —  Preparation  and  Properties  of  Acetylene 
MATERIALS.  —  Calcium  carbide,  acetylene  burner  (for  (c)). 

(a)  Examine  a  typical  specimen  of  calcium  carbide  and  state 
its  characteristic  physical  properties. 

(b)  Fill  a  test  tube  nearly  full  of  water,  stand  it  in  a  rack, 
and  drop  in  two  or  three  very  small  pieces  of  calcium  carbide. 
Acetylene  is  evolved.     After  the  action  has  proceeded  long 
enough  to  expel  the  air,  light  the  gas  by  holding  a  lighted 
match  at  the  mouth  of  the  tube.    Observe  and  record  the  nature 
of  the  flame. 

(c)  Attach  an  acetylene  burner  by  a  short  rubber  tube  to 
a  short  glass  tube  inserted  in  a  one-hole  rubber  stopper.    Put 
10  cubic  centimeters  of  water  in  a  test  tube,  drop  in  a  small 
lump  or  two  of  calcium  carbide,  insert  the  stopper,  and  light 
the  gas  cautiously.    Describe  the  flame. 

SUPPLEMENTARY  EXPERIMENTS 
Experiment  88  —  Properties  of  Graphite 

MATERIALS.  —  Native  and  artificial  graphite. 

(a)  Examine  a  specimen  of  native  and  of  artificial  graphite,  and 
state  the  characteristic  properties  of  each,  especially  the  hardness, 


86 


CHEMISTRY 


cc 


color,  and  luster.      Rub  a  piece  of  graphite  with  the  finger,  and 

describe   the   feeling;    draw  a  piece  slowly  across  a  sheet  of  paper, 

and  state  the  result. 

(b)  If  a  compact  lump  of  graphite  is  available,  determine  its 

specific  gravity  by  the  following  method:  Tie  a  thread  around  the 
solid,  and  weigh  it  on  the  scales  to  a  decigram. 
Slip  it  carefully  into  a  graduated  cylinder  (Fig. 
134)  previously  filled  with  water  to  a  known 
point  and  note  the  increase  in  the  volume  of 
water.  This  increase  in  volume  is  equal  to  the 
volume  of  the  solid.  Calculate  the  specific  grav- 
ity by  dividing  the  weight  of  the  solid  by  the 
weight  of  an  equal  volume  of  water  (assuming 
the  weight  of  i  cubic  centimeter  of  water  to  be  i 
gram).  State  the  result. 

(c)  Wind  the  end  of  a  nichrome  test  wire  (Fig. 
103)   around  a  small  piece  of  graphite  and  hold 
the  graphite  in  the  hottest  part  of  the  flame  for 
a  minute  or  two;  observe  whether  the  graphite 
ignites  readily.      Continue  the  heating  for  five  or 
more  minutes  and  state  the  result. 

(d]  Examine  available  samples  of  graphite  pro- 
ducts, e.g.  stove  polish,  plumbago  crucibles,  core 
of  a  lead  pencil,  electrodes,  lubricants.      Suggest 

a  simple  method  of  testing   them  for  graphite;   verify  it  by  an 
experiment. 

Experiment  89  —  Preparation  of  Carbon  Dioxide  by 
Different  Methods 

(Each  student  need  not  perform  all  of  this  Experiment.} 

MATERIALS.  —  Charcoal,   copper  wire   (30   centimeters  long),    candle, 
magnesium  carbonate,  sodium  carbonate,  sodium  bicarbonate. 

A.  Combustion  of  Carbon.    Wind  one  end  of  the  copper  wire 
around  a  small  lump  of  charcoal,  heat  the  charcoal  in  the  flame,  and 
lower  it  into  a  bottle.    Let  it  remain  for  several  minutes.    Remove  it, 
and  test  the  gas  in  the  bottle  for  carbon  dioxide.    State  the  result. 

B.  Carbonaceous  Substances.     Attach  the  candle  to  the  copper 
wire,  light  the  candle,  and  lower  it  into  a  bottle.    Let  it  burn  a  minute 
or  two,  then  remove  it,  and  test  as  in  A.    State  the  result.    Allow  a 


Fig.  134. — Ap- 
paratus for 
Finding  the 
•S  p  e  c  i  fi  c 
Gravity  of  a 
Solid. 


CARBON  87 

piece  of  wood  and  of  paper  to  burn  in  separate  bottles,  and  test  as  in 
A.     State  the  results. 

C.  Carbonates  and  Acids.     Put  a  little  magnesium  carbonate  into 
a  test  tube,  add  dilute  hydrochloric  acid,  and  test  for  carbon  dioxide 
by  lowering  into  the  escaping  gas  a  tube  which  has  been  dipped  into 
barium  hydroxide  solution.    Observe  and  state  the  change  in  the  drop 
of  barium  hydroxide  solution.    Repeat,  using  sodium  carbonate  and 
dilute  sulphuric  acid;   also  sodium  bicarbonate  and  dilute  sulphuric 
acid.    State  the  results. 

D.  Heating   Carbonates.    Heat  a  little  sodium  bicarbonate  in  a 
test  tube,  and  test  for  carbon  dioxide  as  in  C. 

Experiment  90  —  Carbonic  Acid 

MATERIALS.  —  Solutions    of    sodium    hydroxide  and  phenol-phthalein, 
bottle,  and  a  carbon  dioxide  generator. 

Construct  and  arrange  the  carbon  dioxide  generator  as  in  Exp.  83. 
Fill  a  bottle  half  full  of  water,  add  a  few  drops  of  a  solution  of  phenol- 
phthalein  and  just  enough  sodium  hydroxide  solution  to  color  the 
liquid  a  faint  pink.  Allow  a  slow  current  of  carbon  dioxide  to  bubble 
through  the  liquid  in  the  bottle,  until  a  definite  change  is  produced 
in  the  absorbing  liquid.  Describe  and  explain  it. 

Experiment  91  —  Preparation  and  Properties  of  Carbon 
Monoxide 

MATERIALS.  —  Oxalic    acid,     concentrated     sulphuric     acid,     calcium 
hydroxide  solution,  pneumatic  trough  filled  as  usual,  three  bottles, 
three  glass  plates.    The  apparatus  is  shown  in  Fig.  135. 
Precaution.      Carbon  monoxide  and  oxalic  acid  are  poisonous.       Hot 

sulphuric  acid  is  dangerous.      Perform  this  experiment  with  unusual  care. 

I.  Preparation.  Put  10  grams  of  oxalic  acid  in  the  flask  A,  and 
add  25  cubic  centimeters  of  concentrated  sulphuric  acid.  Put  enough 
calcium  hydroxide  solution  in  B  to  cover  the  end  of  the  tube  E.  The 
end  of  H  should  rest  on  the  bottom  of  the  pneumatic  trough  just 
beneath  the  hole  in  the  support.  Heat  the  flask  A  gently,  and  carbon 
monoxide  will  be  evolved.  A  small  flame  must  be  used,  because  the 
gas  is  rapidly  evolved  as  the  heat  is  increased.  It  is  advisable  to 
remove  or  lower  the  flame  as  bubbles  appear  in  the  tube  B,  —  regu- 
late the  heat  by  the  effervescence.  Collect  all  the  gas,  but  do  not 


88 


CHEMISTRY 


use  the  first  bottle,  covering  the  bottles  with  glass  plates  as  they  are 
filled,  and  setting  them  aside  temporarily.  When  the  last  bottle 
has  been  collected  and  covered,  loosen  the  stopper  in  B,  remove  the 
end  of  H  from  the  water  in  the  trough,  and  if  gas  is  still  being  evolved, 
stand  the  generating  apparatus  in  the  hood.  Proceed  at  once 
with  II. 


Fig.  135.  —  Apparatus  for  Preparing  Carbon  Monoxide. 

II.  Properties.    Test  the  gas  thus:  -(#)  Notice  that  it  is  colorless. 

(£>)  Hold  a  lighted  match  at  the  mouth  of  a  bottle  for  an  instant. 
Note  the  flame,  especially  its  color.  After  the  flame  has  disappeared, 
drop  a  lighted  match  into  the  bottle.  Describe  the  result.  Draw  a 
conclusion  and  verify  it  by  (c). 

(c)  Burn  another  bottle  of  gas,  and  after  the  flame  has  disappeared 
pour  calcium  hydroxide  into  the  bottle  and  shake.  Explain  the 
result. 

Required  Exercises.  —  i.  Describe  briefly  the  preparation  of  carbon 
monoxide. 

2.  Summarize  the  observed  properties  of  carbon  monoxide. 

3.  What  gas  besides  carbon  monoxide  was  produced? 

Experiment  92  —  The  Principle  of  the  Davy  Safety  Lamp 
(a)  Press  a  wire  gauze  down  upon  a  Bunsen  flame.     Where  is 
the  flame?     Remove  the  gauze,  let  it  cool  (or  use  another  gauze), 


CARBON  89 

lower  it  upon  the  flame,  and  hold  a  lighted  match  just  above  the 
gauze.     Now  where  is  the  flame? 

(b)  Extinguish  the  flame.     Turn  on  the  gas,  hold  the  gauze  in 
the  escaping  gas,  about  15  centimeters  (6  inches)  above  the  top  of 
the  burner,  and  thrust  a  lighted  match  into  the  gas  above  the  gauze. 
Where  is  the  flame?    Lower  the  gauze  slowly  and  describe  the  final 
result. 

(c)  Hold  the  gauze  in  the  flame  in  one  position  for  a  minute  or 
two.    Where  is  the  flame  at  the  end  of  this  time?    Why? 

Required  Exercises.  —  i.    Define  kindling  temperature. 

2.    State  exactly  how  this  experiment  illustrates  kindling  temperature. 

Experiment  93  —  Properties  of  Carborundum 

Examine  specimens  of  different  varieties  of  carborundum  and  state 
the  characteristic  properties  of  each,  especially  the  hardness. 


ILLUMINATING   GAS  — FLAME 

Experiment  94  —  Preparation  and  Properties  of 
Illuminating  (Coal)  Gas 

MATERIALS.  —  Soft  coal,  litmus  paper,  filter  paper,  lead  nitrate  solution. 
Arrange  an  apparatus  like  the  A-B  part  shown  in  Fig.  108. 
Fill  the  large  test  tube  A  two-thirds  full  of  coarsely  powdered 
soft  coal,  insert  the  stopper  with  its  delivery  tube  B,  and 
clamp  the  test  tube  carefully  to  the  iron  stand  as  shown  in 
Fig.  108.  Heat  the  whole  tube  gently  at  first,  and  gradually 
increase  the  heat,  but  avoid  heating  either  end  very  hot. 

(a)  As  soon  as  the  gas  begins  to  escape,  hold  at  the  end 
of  the  tube  B  a  piece  of  filter  paper  which  has  been  moistened 
with  lead  nitrate  solution;   observe  the  effect  upon  the  paper. 
The  discoloration  is  caused  by  lead  sulphide  which  is  produced 
by  the  interaction  of  lead  nitrate  and  the  sulphides  in  the 
liberated  gas. 

(b)  Lay  a  piece  of  moistened  red  litmus  paper  on  the  end  of 
the  tube  B  and  continue  to  heat  strongly.    Observe  any  change 
in  the  litmus  paper.     To  what  compounds  in  the  gas  is  the 
change  due? 

(c)  Heat  strongly,  and  light  the  gas  at  the  end  of  the  tube 
B.     Observe  and  describe  the  flame. 

(d)  Discontinue  heating,  let  the  apparatus  cool  somewhat, 
disconnect,  and  break  open  the  test  tube.    Examine  the  con- 
tents.   State  the  properties  of  both  solid  and  liquid  products; 
what  is  the  name  of  each? 

Experiment  95  —  Candle  Flame 

MATERIALS.  —  Candle,  two  blocks  of  wood,  bottle,  piece  of  stiff  white 
paper,  calcium  hydroxide  solution,  matches,  a  lead  pencil,  copper 
wire  (15  centimeters  or  6  inches  long). 

Attach  a  candle  to  a  block  of  wood  by  means  of  a  little 
melted  candle  wax,  and  proceed  as  follows: 


ILLUMINATING   GAS  —  FLAME  91 

(a)  Hold  a  cold,  dry  bottle  over  the  lighted  candle.    Describe 
the  result  produced  inside  the  bottle.     What  is  the  product? 
What  is  its  source?    Remove  the  bottle,  pour  a  little  calcium 
hydroxide  solution  into  it,  and  shake.  Describe  and  explain  the 
result.     What  are  the  two  main  products  of  a  burning  candle? 

(b)  Blow  out  the  candle  flame,  and  immediately  hold  a 
lighted  match  in  the  escaping  smoke.     Does  the  candle  re- 
light?    Why?     What  is  the  general  nature  of  this  smoke? 
How  is  it  related  to  the 

candle  wax?  How  does 
(b)  contribute  to  the 
explanation  of  (a)? 

(c)  Press  a  piece   of 
stiff  white  paper  for  an 
instant  down  upon  the 

candle  flame  almost  to  FiS'  ^6-  ~  Effect  of  CoolinS  a 

.  ,         -„  Candle  Flame, 

the  wick.     Repeat  sev- 
eral times  with  different  parts  of  the  paper.      What  does  the 
paper  show  about  the  structure  of  the  flame? 

(d)  Roll  one  end  of  the  copper  wire  around  a  lead  pencil 
to  form  a  spiral  about  (2  centimeters  or  i  inch)  long.     Press 
the  spiral  down    slowly  upon   the  candle  flame   (Fig.    136). 
Repeat  after  cooling  the  wire.   What  is  the  result?  Why? 

Optional  Exercises.  —  i.  Draw  a  candle  flame,  showing  the  parts. 

2.  What  is  the  essential  difference  between  a  candle  flame  and  a 
Bunsen  burner  flame? 

3.  Is  there  any  essential  difference  between  a  candle  and  a  gas  or  a 
lamp  flame? 

4.  Why  do  candles  and  lamps  often  smoke? 

Experiment  96  —  Bunsen  Burner  and  Bunsen  Burner 
Flame 

MATERIALS.  —  Bunsen   burner,   glass   tube,   powdered   charcoal,   pin 

copper  wire. 

A.  Bunsen  Burner.  Take  apart  a  Bunsen  burner  and 
study  the  construction.  Write  a  short  description  of  the 
burner.  Sketch  the  essential  parts. 


92  CHEMISTRY 

B.  Bunsen  Burner  Flame,  (a)  Close  the  holes  at  the  bot- 
tom of  a  lighted  burner  and  hold  a  glass  tube  in  the  upper 
part  of  the  yellow  flame.  Note  the  black  deposit.  What  is 
it?  Where  did  it  come  from?  Open  the  holes  and  move  the 
blackened  tube  up  and  down  in  the  colorless  Bunsen  flame. 
What  becomes  of  the  deposit? 

(b)  Dip  a  glass  tube  a  short  distance  into  some  powdered 
wood  charcoal,  place  the  end  containing  the  charcoal  in  one 
of  the  holes  at  the  bottom  of  the  lighted  burner,  and  blow 
gently  two  or  three  times  into  the  other  end.     Describe  and 
explain  the  result. 

(c)  Open  and  close  the  holes  of  a  lighted  burner  several 
times.     Describe  the  result.    Pinch  the  rubber  tube  to   ex- 
tinguish the  flame,  then  light  the  gas  at  the  holes.     What 
change  is  produced  in  the  flame?     What  is  the  object  of  the 
holes? 

(d)  Hold  a  match  across  the  top  of  the  tube  of  a  lighted 
Bunsen  burner.    When  the  match  begins  to  burn,  remove  and 
extinguish  it.    Note  where  it  is  charred,  and  explain  the  result. 
Press  a  piece  of  wire  gauze  down  upon  the  flame.     Describe 

the  appearance  of   the  gauze.     The  same 
w  fact    may  be   shown   by    sticking     a    pin 

J  L  e        through  a  (sulphur)   match,  suspending  it 

across  the  burner,    and  then  lighting   the 
gas.    The  position  of  the  match  is  shown  in 
Fig.  137.    Turn  on  a  full  current  of  gas  be- 
fore lighting  it.     What  does  the  whole  of 
experiment  (d)  show   about   the   structure 
Fig.  137.  — Sulphur  of  the  lower   part   of   the   Bunsen   flame? 
Match  Suspended    Verif  angwer  b     (^ 

Across  the  Top  of         /  \   TT  u  j     *  i_      /  i_ 

a  Bunsen  Burner.        W   H°ld  One  end  °f  a  SlaSS  tube    fabout 
15  centimeters  or   6    inches    long)    in  the 

Bunsen  flame  about  2  centimeters  (i  inch)  from  the  top  of 
the  burner  tube.  Hold  a  lighted  match  for  an  instant  at  the 
upper  end  of  the  tube;  raise  or  lower  the  tube  slightly  (still 
keeping  the  end  in  the  flame)  and  observe  the  result.  What 


ILLUMINATING   GAS  —  FLAME  93 

does   the   result   show   about    the    structure   of  the    Bunsen 
flame?    How  does  it  verify  (d)? 

(/)  Find  the  hottest  part  of  the  flame,  when  a  full  current 
of  gas  is  burning,  by  holding  a  copper  wire  in  the  flame.  Meas- 
ure its  distance,  approximately,  from  the  top  of  the  burner 
tube. 

(g)  Examine  a  slightly  imperfect  Bunsen  burner  flame  — 
one  which  shows  clearly  the  outlines  of  the  inner  part.    What 
is  the  general  shape  of  each  main  part?    Draw  a  vertical  and 
a  cross  section  of  the  flame. 

Experiment  97  —  Reduction  and  Oxidation  with  the 
Blowpipe 

(Each  pupil  need  not  perform  all  of  this  Experiment.) 

MATERIALS.  —  Blowpipe,  blowpipe  tube,  charcoal,  lead  oxide,  sodium 
carbonate,  sodium  sulphate,  wood  charcoal,  silver  coin,  zinc, 
lead,  tin. 

Slip  the  blowpipe  tube  into  the  burner.    Light  the  gas  and 
lower  the  flame   until   it 
is    about    4    centimeters      /^ 
(1.5  inches)   high.     Rest     ([ 

the  tip  of  the  blow  pipe  Fig>  I38>  _  Blowpipe. 

(Fig.  138)  on  the  top  of 

the  tube,  placing  the  tip  just  within  the  flame.    Put  the  other 
end  of  the  blowpipe  between  the  lips,  puff  out  the  cheeks, 
inhale  through  the  nose,  and  exhale  into  the  blowpipe,  using 
the  cheeks  somewhat  as  bellows.     Do  not 
blow  in  puffs,  but  produce  a  continuous 
flow  of  air  by   steady  and  easy  inhaling 
Fig.  139. —  Blowpipe   and  exhaling.      The  operation  is  natural 
Flame    A  (oxidiz-  ancj   sjmpiej  and,   if   properly  performed, 
'  will  not   make   one   out   of  breath.     The 
flame  should  be  an  inner  blue  cone  sur- 
rounded by  an  outer  and  almost  invisible  cone,  though  its  shape 
varies  with  the  method  of  production  (Fig.  139).     Practice  un- 
til the  flame  is  produced  voluntarily  and  without  exhaustion. 


94  CHEMISTRY 

A.  Reduction,     (a)  Make  a  shallow  hole  at  one  end  of  the 
flat  side  of  a  piece  of  charcoal.    Fill  the 'hole  with  a  mixture 
of  equal  parts  of  powdered  sodium  carbonate  and  lead  oxide, 
and  heat  the  mixture  in  the  reducing  flame.    In  a  short  time 
bright,  silvery  globules  will  appear  on  the  charcoal.    Let  the 
mass  cool,  and  pick  out  the  largest  globules.    Put  one  or  two 
in  a  mortar,  and  strike  with  a  pestle.    Are  they  soft  or  hard? 
Malleable  or  brittle?     How  do  the  properties  compare  with 
those  of  metallic  lead?    What  has  become  of  the  oxygen? 

(b)  Grind  together  in  a  mortar  a  little  sodium  sulphate  and 
wood  charcoal,  adding  at  intervals  just  enough  water  to  hold 
the  mass  together.  Heat  some  of  this  paste  in  the  reducing 
flame  as  in  (a).  Scrape  the  fused  mass  into  a  test  tube,  boil  in 
a  little  water,  and  put  a  drop  of  the  solution  on  a  bright  silver 
coin.  If  a  dark  brown  stain  is  produced,  it  is  evidence  of  the 
formation  of  silver  sulphide.  Repeat,  if  no  such  stain  is  pro- 
duced. The  silver  sulphide  is  formed  by  the  interaction  of 
silver  and  sodium  sulphide.  Explain  how  the  experiment 
illustrates  reduction. 

B.  Oxidation,     (a)  Heat  a  small  piece  of  zinc  on  charcoal 
in  the  oxidizing  flame.    Direct  the  flame  so  that  most  of  the 
product  will  form  a  coating  on  the  charcoal.     What  is  the 
product?     Observe  the  color  of  the  coating  on  the  charcoal 
when  hot  and  cold.    Record  the  result. 

(b)  Heat  a  piece  of  lead  as  in  (a).     Observe  the  color  of 
the  coating  when  hot  and  cold.     Record  the  result. 

(c)  Heat  a  small  piece  of  tin  in  the  oxidizing  flame.    Observe 
and  record  as  in  (b). 

Optional  Exercises.  —  i.    Name  the  products  formed  in  B. 

2.  Sketch  a  blowpipe. 

3.  Sketch  a  flame  showing  the  oxidizing  and  reducing  parts. 

SUPPLEMENTARY  EXPERIMENTS 
Experiment  98  —  Combustion  of  Illuminating  Gas 

MATERIALS.  —  Pointed  glass  tube,  calcium  hydroxide  solution,  bottle. 

Remove  the  Bunsen  burner  from  the  rubber  connection  tube  and 

replace  it  by  a  glass  tube  with  a  small  opening.    Light  the  gas,  and 


ILLUMINATING   GAS  —  FLAME  95 

lower  a  small  flame  into  a  cold,  dry  bottle.  Observe  at  once  the  most 
definite  result  inside  the  bottle.  Remove  and  extinguish  the  flame, 
pour  a  little  calcium  hydroxide  solution  into  the  bottle,  and  shake. 
What  are  the  two  products  of  the  combustion  of  illuminating  gas? 

Experiment  99  —  Properties  of  the  By-Products  of  the 
Manufacture  of  Illuminating  Gas 

MATERIALS.  —  Tar,  ammoniacal  liquor,  coke,  gas  carbon. 

A.  Tar.    Examine  a  specimen  of  tar  and  state  its  characteristic 
properties. 

B.  Ammoniacal  Liquor,     (a)  Proceed  as  in  A,  using  ammoniacal 
liquor. 

(b)  Recall,  perform,  or  repeat  (if  necessary)  Exp.  53  C. 

C.  Coke  and  Gas  Carbon,    (a)  Proceed  as  in  A,  using  coke  and  gas 
carbon. 

(b)  Proceed  with  coke  and  with  gas  carbon  as  in  Exp.  88. 

(c)  Compare  with  the  results  obtained  in  Exp.  88  (c). 

Experiment  100  —  Testing  Illuminating  Gas 

(a)  Test  samples  of  illuminating  gas  for  carbon  dioxide,  sulphides 
(Exp.  94  (a)  ),  and  ammonia.    State  the  results. 

(b)  Suggest  a  test  for  carbon  monoxide.    Submit  the  details  to  the 
Teacher,  before  proceeding. 

(c)  Test  illuminating  gas  for  moisture. 

Experiment  101  —  Testing  Metals  with  the  Blowpipe 

Obtain  "unknowns"  from  the  Teacher  and  test  as  in  Exp.  97  B. 

Experiment  102  —  Welsbach  Burner,  Mantle,  and 
Flame 

MATERIALS.  —  Welsbach  burner  and  mantle. 

Examine  a  Welsbach  burner  and  compare  its  structure  with  that 
of  the  Bunsen  burner.  Connect  the  burner  with  the  gas  supply, 
light  the  gas,  and  compare  the  flame  with  the  Bunsen  burner  flame. 
Are  the  burners  and  flames  essentially  different? 

Examine  a  Welsbach  mantle  carefully.  Suspend  the  mantle  on 
the  end  of  an  iron  wire  (e.g.  the  handle  of  a  deflagrating  spoon),  and 
hold  the  mantle  in  the  Bunsen  burner  flame.  State  the  result. 


ORGANIC   COMPOUNDS  — FOOD 

Experiment  103  —  Composition  of  Organic  Compounds 

MATERIALS.  —  Meat,  gelatin,  glue,  leather,  albumin,  mustard,  sugar. 

A.  Carbon,  (a)  Recall,   perform,   or   repeat   (if   necessary) 
the  experiment  showing  the  distribution  of  carbon  (Exp.  80). 

(b)  Heat  a  very  small  piece  of  meat  in  a  test  tube  and 
state  the  final  result.     (See  also  Exp.  104  (d)  and  117  A.) 

B.  Nitrogen.     Proceed  as  in   Exp.   117    B,   using  various 
organic    substances,    e.g.  gelatin,  glue,   meat,   leather,   peas, 
beans,  nuts,  leaves.     State  each  result. 

C.  Sulphur.     Proceed  as  in  Exp.  117  C,  using  various  or- 
ganic substances,  e.g.  albumin,  mustard.      State  each  result. 

D.  Phosphorus.     Proceed  as  in  Exp.  117  D,  using  various 
organic  substances,  e.g.  albumin,  casein,  seeds,   cereals,  nuts. 
State  each  result. 

E.  Hydrogen  and  Oxygen.    Heat  a  little  dry  sugar  in  a  test 
tube  and  notice  the  deposit  of  water  in  the  upper  part  of  the 
test  tube. 

Experiment  104  —  Properties  of  Sucrose  (Cane  Sugar) 
MATERIAL.  —  Cane  sugar. 

(a)  Examine  different  varieties  of  sucrose  and  state  the 
characteristic  properties. 

(b)  Put  a  little  sucrose  in  the  upper  end  of  a  test  tube, 
hold  the  test  tube  in  a  horizontal  position,  and  heat  very 
gently  by  moving  it  back  and  forth  above  the  flame.    As  soon 
as  the  sugar  is  melted,  pour  a  little  out  upon  a  glass  plate. 
Examine  it  later  and  describe  the  cooled  mass.     Add  more 
sugar,  heat  as  usual,  and  notice  the  change  in  color  of  the 
substance.     Note  the  odor  and  also  the  deposit  of  water  in 
the  upper  part  of  the  test  tube.    Heat  still  further,  and  describe 


ORGANIC  COMPOUNDS  —  FOOD      97 

the  substance  finally  obtained.  When  the  test  tube  is  cool, 
break  it  and  examine  the  residue.  What  is  it?  Verify  your 
answer  by  a  simple  test. 

(c)  If  time  permits,  prepare  crystals  of  sucrose  by  sus- 
pending a  thread  in  a  concentrated  solution.     Describe  the 
crystals. 

(d)  Put  a  little  cane  sugar  in  a  test  tube,  add  enough  con- 
centrated sulphuric  acid  to  cover  it,  and  mix  by  shaking. 
Observe  the  result  after  a  few  minutes.    If  the  change  is  not 
conspicuous,  warm  slightly.     What  in  all  probability  is  the 
final  solid  substance? 

Experiment  105  —  Properties  of  Dextrose  (Glucose) 

MATERIALS.  —  Glucose,  solutions  of  silver  nitrate  and  sodium  hydroxide. 

(a)  Examine  different  varieties  of  dextrose,  e.g.  glucose, 
grape  sugar,  and  state  the  characteristic  properties.  Taste 
and  compare  the  sweetness  with  that  of  sucrose. 

(6)  Proceed  as  in  Exp.  104  (b),  using  grape  sugar  instead  of 
sucrose.  Compare  the  results,  especially  the  color  and  odor. 

(c)  Put  a  little  grape  sugar  in  a  test  tube,  add  concentrated 
sulphuric  acid,  and  examine  after  a  short  time.    Compare  with 
Exp.  104  (d). 

(d)  For  fermentation,  see  Exp.  127. 

(e)  For  Fehling's  test,  see  Exp.  106. 

(/)  Clean  a  test  tube  thoroughly  by  boiling  dilute  sodium 
hydroxide  solution  in  it,  and  washing  several  times  with  water. 
Put  10  cubic  centimeters  of  silver  nitrate  solution  in  the  clean 
test  tube,  and  add  ammonium  hydroxide  slowly  until  the  pre- 
cipitate at  first  formed  redissolves,  taking  care  to  mix  the 
solutions.  Add  a  little  grape  sugar  solution  and  warm  gently. 
Silver  will  be  deposited  as  a  bright  film  inside  the  test  tube. 

(e)  If  a  polafiscope  is  available,  examine  a  solution  of  dex- 
trose according  to  directions  given  by  the  Teacher. 


98  CHEMISTRY 

Experiment  106  —  Fehling's  Test  for  Sugar 

MATERIALS.  —  Copper  sulphate,  Rochelle  salt,  sodium  hydroxide,  and 
sugar  solutions. 

(a)  Mix  equal  (and  small)  volumes  of  copper  sulphate  and 
Rochelle  salt  solutions  in  a  test  tube,  and  boil  carefully;   then 
add  enough  sodium  hydroxide  solution  to  make  the  mixture 
strongly  alkaline.1     Add  a  little  glucose  solution,   and  boil 
until  a  decided  change  is  produced.    The  precipitate  is  cuprous 
oxide.     Describe  it. 

(b)  Repeat  (a),  using  cane  sugar  solution  instead  of  glu- 
cose.    State  the  result. 

Experiment  107  —  Testing  for  Glucose 

MATERIALS.  —  Fehling's  solution  and  the  substances  enumerated  below. 
Apply  Fehling's  test  for  glucose  (and  similar  sugars)  to 
cheap  candy,  maple  sugar,  molasses,  table  sirups,  jelly,  jam, 
etc.  Prepare  and  use  clear  solutions.  State  the  result  in  each 
case. 

Experiment  108  —  Properties  of  Starch 

MATERIALS.  —  Starch,  microscope,  Fehling's  solution,  iodine  solution. 

(a)  Examine  different  kinds  of  starch  with  a  microscope, 
if  one  is  available.     Describe  them.     If  time  permits,  make 
drawings.    . 

(b)  Examine  with  a  microscope  thin  slices  of  peas  or  beans 
which  have  been  soaked  about  eight  hours  in  water.    Describe. 

(c)  Examine   a   thin   slice   of  potato   with   a   microscope. 
Describe. 

(d)  Prepare  a  starch  mixture  by  boiling  about   i   gram  of 
powdered  starch  for  a  few  minutes  in  a  test  tube  containing 

1  This  mixture,  which  is  called  Fehling's  solution,  may  be  prepared 
accurately  as  follows:  Dissolve  34.64  gm.  of  crystallized  copper  sulphate 
in  500  cc.  of  water  —  solution  No.  i;  dissolve  180  gm.  of  Rochelle  salt 
(sodium  potassium  tartrate)  and  70  gm.  of  sodium  hydroxide  in  500 
cc.  of  water — solution  No.  2.  Filter,  if  not  clear.  Mix  the  two  solu- 
tion (equal  volumes)  just  before  using. 


ORGANIC   COMPOUNDS  —  FOOD  99 

50  cubic  centimeters  of  water;  stir  or  agitate  the  mixture  dur- 
ing the  boiling.  Make  three  tests  with  the  starch  mixture  (the 
third  is  (e)  below),  (i)  Pour  most  of  it  into  an  evaporating 
dish  which  stands  on  a  gauze-covered  ring  and  boil,  add  i 
cubic  centimeter  of  concentrated  sulphuric  acid,  mix  well,  and 
boil  for  at  least  ten  minutes;  add  water  occasionally  to  replace 
that  lost  by  evaporation.  Meanwhile  proceed  with  the  second 
test.  (2)  Dilute  the  rest  of  the  original  starch  mixture  with 
water  and  test  it  with  Fehling's  solution.  Observe  and  state 
the  final  result.  As  soon  as  the  mixture  in  (i)  has  been  boiled 
at  least  ten  minutes,  take  out  a  little,  add  sodium  hydroxide 
solution  to  alkaline  reaction  and  apply  Fehling's  test.  Note 
the  result.  Continue  the  heating  for  ten  or  more  minutes,  and 
test  again.  State  the  final  result. 

(e)  Prepare  a  dilute  iodine  solution  by  dissolving  a  few 
crystals  of  iodine  in  10  cubic  centimeters  of  alcohol.  Add  a 
few  drops  of  the  iodine  solution  to  a  dilute,  cold  starch  mix- 
ture. Observe  the  blue  color.  (This  test  for  starch  is  delicate, 
and  dilute  mixtures  should  be  used.) 

Experiment  109  —  Properties  of  Alcohol 

MATERIALS.  —  Alcohol,  camphor,  shellac,  rosin,  porcelain  dish. 

A.  Ethyl  Alcohol,     (a)  Determine  cautiously  the  odor  and 
taste  of  alcohol.    Drop  a  little  on  a  glass  plate  or  on  a  piece 
of  paper,  and  watch  it  evaporate.    Is  its  rate  of  evaporation 
more  rapid  than  that  of  water? 

(b)  Weigh  a  measured  quantity  (about  25  cubic  centimeters) 
of  alcohol  and  calculate  its  specific  gravity. 

(c)  Alcohol  dissolves  many  organic  substances.     Try  cam- 
phor, powdered  shellac,  or  rosin.    Describe  the  result.    Verify 
the  solvent  power  of  alcohol  by  adding  water  to  the  solutions. 
Describe  the  result. 

(d)  Burn  a  little  alcohol  in  a  dish  and  observe  the  nature  of 
the  flame.    What  are  the  products  of  combustion? 

B.  Methyl  Alcohol.     Repeat  A,  using  methyl  alcohol. 


ioo  CHEMISTRY 

Experiment  110  —  Properties  of  Acetic  Acid 
Treat  acetic  acid  as  follows:    (a)  Taste  (cautiously),  and 
describe. 

(b)  Test  with  litmus  paper,  and  describe  the  result. 

(c)  Warm  a  little  in  a  test  tube,  and  smell  (cautiously). 
Describe  the  odor. 

Experiment  111  —  Test  for  Acetic  Acid  and  Acetates 

Cautiously  add  a  few  drops  of  concentrated  sulphuric  acid 
to  equal  (and  small)  volumes  of  acetic  acid  and  ethyl  alcohol. 
Shake  the  mixture  and  warm  gently.  The  pleasant,  fruitlike 
odor  is  due  to  ethyl  acetate. 

NOTE.  —  This  experiment  is  also  a  test  for  alcohol. 

Experiment  112  —  Properties  of  Vinegar 

(a)  Show,  experimentally,  that  a  sample  of  vinegar  contains 
acetic  acid. 

(b)  Evaporate  a  little  vinegar  to  dryness  on  a  water  bath 
and  note  the  residue.    Stand  the  dish  on  a  gauze-covered  ring 
and  heat  gently.    Note  the  ash.    Test  the  ash  for  (i)  a  car- 
bonate and  (2)  potassium.     Perform  (2)  by  heating  a  little 
of  the  ash  with  a  test  wire;   potassium  compounds  color  the 
flame  a  delicate  lilac. 

Experiment  113  —  Testing  Baking  Powder 

MATERIALS.  —  Baking  powder,  vinegar,  sour  milk,  lemon  juice,  iodine 
solution,  sodium  hydroxide,  ammonium  oxalate  solution,  and  the 
substances  needed  for  B,  C,  E,  G. 

A.  Carbonates,  (a)  Put  a  little  baking  powder  in  a  test 
tube,  add  a  few  drops  of  dilute  hydrochloric  acid,  and  test  the 
escaping  gas  with  a  tube  which  has  been  dipped  into  barium 
hydroxide  solution.  State  the  result. 

(b)  Put  a  little  baking  powder  in  a  test  tube,  add  15  to  20 
cubic  centimeters  of  water,  and  shake  well.  Let  the  action 
continue  a  short  time,  and  then  test  as  in  (a).  State  the  result. 


ORGANIC   COMPOUNDS  —  FOOD  101 

(c)   Add  to  a  little  baking  powder  sour  substances,   e.g. 
vinegar,  sour  milk,  lemon  juice,  and  state  the  result. 

B.  Tartrates.     Prepare  a  cold  solution  of  baking  powder  by 
shaking  about  3  grams  of  the  substance  with  25  cubic  centi- 
meters of  water,  filter,  if  not  clear,  and  use  the  clear  solution 
in  this  and  succeeding  experiments  (except  D  and  G).     Clean 
a  test  tube  by  boiling  sodium  hydroxide  solution  in  it  and 
then  washing  thoroughly  with  water,  and  then  proceed  as  in 
Exp.  105   (/),  using  the  baking  powder  solution  instead   of 
glucose.     Tartrates,  if  present,   will  reduce  the  silver  com- 
pound to  silver,  which  will  coat  the  inside  of  the  test  tube. 

C.  Chlorides  and  Sulphates.    Apply  the  usual  tests  for  these 
to  small  portions  of  the  baking  powder  solution  (prepared  in 
B).    In  each  case  acid  must  be  added  to  acid  reaction  and  the 
solution  boiled  before  testing  —  dilute  nitric  for  chlorides  and 
dilute  hydrochloric  for  sulphates.    State  the  result. 

D.  Starch.     Apply  the  iodine  test  for  starch   to   a   little 
baking  powder  mixed  with  water.    (See  Exp.  108  (e) ).     State 
the  result. 

E.  Phosphates.     Warm  a  little  of  the  baking  powder  solu- 
tion with  a  little  concentrated  nitric  acid,  and  test  as  in  Exp. 
117  D.    State  the  result. 

F.  Ammonium  Compounds.     Boil  a  few  cubic  centimeters 
of  the  baking  powder  solution  with  an  equal  volume  of  sodium 
hydroxide  solution.     The  presence  of  ammonium  compounds 
is  shown  by  the  liberation   of   ammonia  gas,  which  can  be 
detected   by  its  odor  and  its  action  on  red  litmus  paper. 
State  the  result. 

G.  Aluminium  Compounds,    (i)  Proceed  as  in  Exp.  209  (c) 
using  a  little  baking  powder  instead  of  an  aluminium  com- 
pound.     (2)  Heat  some  baking  powder  in  a  porcelain  dish. 
Add  hot  water,  boil,  filter,  and  add  considerable  ammonium 
chloride  to  the  filtrate.      A  whitish  flocculent  precipitate  of 
aluminium  hydroxide  is  produced.      State  the  result. 

H.   Calcium  Compounds.     Boil  a  few  cubic  centimeters  of 
the  clear  baking  powder  solution  with  dilute  hydrochloric 


102  CHEMISTRY 

acid  (to  remove  the  carbon  dioxide),  add  ammonium  hydroxide 
to  alkaline  reaction,  filter  if  not  clear,  and  then  ammonium 
oxalate  solution.  If  calcium  compounds  are  present,  a  white 
precipitate  (calcium  oxalate)  will  be  formed. 

Experiment  114  —  General  Properties  of  Fats 

MATERIALS.  —  Fats,  liquids  as  in  (c). 

(a)  Examine  several  kinds  of  fat,  e.g.  lard,  tallow,  butter, 
and  note  the  consistency,  color,  odor,  taste,  and  any  other 
characteristic  property. 

(b)  Put  a  small  quantity  of  fat  on  a  piece  of  filter  paper, 
drop  a  little  alcohol  or  ether  on  another  part  of  the  same 
paper,  and  compare  the  results  after  a  few  minutes. 

(c)  Are  fats  soluble  in  water,  alcohol,  ether,  gasoline,  and 
carbon  tetrachloride?     Try  the  effect  of  one  or  more  of  these 
liquids  (e.g.   gasoline)  separately  on  several  fats.    State  the 
result. 

(d)  Test  rancid  fat  for  acid  by  shaking  the  fat  with  a  little 
alcohol  and  then  adding  a  few  drops  of  neutral  litmus  or 
phenol-phthalein  solution  (or  a  piece  of  sensitive  litmus  paper) 
to  the  solution.     State  the  result. 

Experiment  115  —  Preparation  of  Soap 

MATERIALS.  —  Sodium  hydroxide,  lard,  salt;  for  C  potassium  hydroxide, 
alcohol,  cotton-seed  oil. 

Prepare  soap  by  one  of  the  following  methods :  A.  Dissolve 
10  grams  of  sodium  hydroxide  in  75  cubic  centimeters  of  water, 
add  30  grams  of  lard,  and  boil  the  mixture  in  a  metal  dish  for 
an  hour  or  more;  add  water  occasionally  to  replace  that  lost 
by  evaporation.  Then  add  20  grams  of  fine  salt  in  small 
portions.  Stir  constantly  during  the  addition  of  the  salt.  Let 
the  mass  cool,  and  then  remove  the  soap,  which  will  form  in 
a  cake  at  the  surface. 

B.  Dissolve  13  to  15  grams  of  sodium  hydroxide  in  100  cubic 
centimeters  of  water,  add  100  cubic  centimeters  of  castor  oil, 


ORGANIC   COMPOUNDS  —  FOOD  103 

and  boil  for  about  half  an  hour.    Add  20  grams  of  salt,  and 
then  proceed  as  in  A. 

C.  Dissolve  i  gram  of  potassium  hydroxide  in  10  cubic 
centimeters  of  alcohol,  add  a  little  lard  or  cotton-seed  oil, 
and  stir  constantly  while  the  mixture  is  being  heated  cautiously 
to  sirupy  consistency.  Allow  the  solution  to  cool.  The  jelly- 
like  product  is  soap. 

Experiment  116  —  Properties  of   Soap 

MATERIALS.  —  Soap,  sulphuric  acid,  calcium  sulphate,  magnesium  sul- 
phate, and  acid  calcium  carbonate  solutions. 

Test  as  below  the  soap  prepared  in  Exp.  115  (or  another 
sample,  if  desired). 

(a)  Leave   soap   shavings   exposed   to   the   air  for   several 
days.     What  does  the  result   show  about  the  presence    of 
water  in  the  soap? 

(b)  Test    soap    solution    with    litmus    paper.       State    the 
result.     Put  a  few  drops  of  phenol-phthalein  solution  on  dry 
soap.      State  the  result.      This  is  a  test  for  "  free  alkali." 

(c)  Add  considerable  dilute  sulphuric  acid  to  a  soap  solu- 
tion.    The  precipitate  is  a  mixture  mainly  of  palmitic  and 
stearic  acids.     Describe  it. 

(d)  To  a  little  soap  solution  in  separate  test  tubes  add  cal- 
cium sulphate  and  magnesium  sulphate  solutions.     Describe 
the  result.     Boil  for  a  few  minutes  and  describe  the  result. 
Prepare   a   solution   of   acid   calcium    carbonate   by   passing 
carbon  dioxide  into  limewater  until  the  precipitate  is  redis- 
solved  (see  Exp.  85).     Add  some  of  the  solution  to  a  soap 
solution,  and  describe  the  result.    Boil,  as  above,  and  describe 
the  result. 

Experiment  117  —  Composition  of  Proteins 

MATERIALS.  —  Albumin,  soda-lime,  sodium  hydroxide  solution,  lead 
acetate  solution,  sodium  carbonate,  potassium  nitrate,  ammonium 
molybdate  solution. 

A.  Carbon.  Put  a  few  small  pieces  of  dry  egg  albumin  in 
a  test  tube.  Heat  gently  at  first  and  finally  increase  the  heat. 


104  CHEMISTRY 

Notice  the  characteristic  odor  of  the  burning  albumin.    Notice 
also  the  change  in  color. 

B.  Nitrogen.     Put  a  small  amount  of  dry  egg  albumin  in 
a  test  tube  and  add  five  times  its  bulk  of  soda-lime.    Mix  by 
shaking.    Heat  gently  and  test  the  escaping  vapors  with  moist 
red  litmus  paper.     Explain  the  result. 

C.  Sulphur.     Put  a  small  quantity  of  dry  egg  albumin  in 
a  test   tube.    Add   5   to   10  cc.  of  dilute  sodium  hydroxide 
solution  and  a  few  drops  of  lead  acetate  solution.     Boil  the 
mixture  a  few  moments  and  notice  the  change  in  color.    The 
black  (or  brown)  compound  is  lead  sulphide. 

D.  Sulphur  and  Phosphorus.     Put  a  small  quantity  of  dry 
egg  albumin  in  a  porcelain  dish  and  add  about  five  times  its 
bulk  of  fusion  mixture,  i.e.  equal  quantities  of  powdered  sodium 
carbonate  and  potassium  nitrate.    Stand  the  dish  on  a  gauze- 
covered  ring  attached  to  an  iron  stand.     Heat  cautiously  at 
first  and  finally  increase  the  heat  sufficiently  to  produce  mild 
deflagration.     Continue  the  heat  until  the  mixture  becomes 
nearly  colorless.     Let  the  dish  cool.     Meanwhile  boil  some 
water  in  a  test  tube  and  dissolve  the  residue  in  the  boiling  water. 
Filter  the  solution  if  it  is  not  clear,  and  divide  into  two  parts. 
To  one  part  add  concentrated  hydrochloric  acid  drop  by  drop, 
boil  until  effervescence  ceases,  and  then  add  barium  chloride 
solution.    What  is  the  white  precipitate?    Explain  its  forma- 
tion.   To  the  other  portion  add  concentrated  nitric  acid  drop 
by  drop,  boil  until  effervescence  ceases,  and  then  add  a  little 
ammonium  molybdate  solution.      Warm   slightly.      A  yellow 
precipitate    of    ammonium    phospho-molybdate    should    be 
formed. 

Experiment  118  —  Tests  for  Proteins 

MATERIALS.  —  Albumin  and  the  specified  solutions. 

A.  Color  Reactions,  (a)  Prepare  a  dilute  solution  of  egg 
albumin.  To  about  5  cc.  add  an  equal  volume  of  sodium 
hydroxide  solution.  Then  add  drop  by  drop  a  dilute  copper 
sulphate  solution.  A  violet  color  is  produced. 


ORGANIC   COMPOUNDS  —  FOOD  105 

(b)  To  5  cc.  of  albumin  solution  add  an  equal  volume  of 
concentrated  nitric  acid.  Heat  gently  until  a  yellow  precipi- 
tate or  a  yellow  solution  is  obtained.  Cool  in  running  water 
and  add  an  excess  of  sodium  hydroxide  solution.  An  orange 
color  is  produced. 

B.  Precipitation,     (a)  To   5   cc.   of  albumin  solution   add 
concentrated  nitric  acid  slowly,  pouring  the  acid  down  the 
inside  of  the  tube  so  the  two  solutions  will  not  mix.    A  white 
cloudy  precipitate  is  formed  at  the  surface  of  the  two  liquids. 

(b)  Add  a  few  drops  of  mercuric  chloride  solution  (POISON) 
to  a  little  albumin  solution.     Observe  the  white  precipitate. 

(c)  Proceed  as  in  (b),  using  alum  solution  instead  of  mer- 
curic chloride  solution. 

C.  Coagulation.     Put  about  5  cc.  of  undiluted  egg  albumin 
in  a  test  tube  and  heat  gently.    Observe  and  state  the  result. 

Experiment  119  —  Testing  Food 
MATERIALS.  —  Samples  of  food  from  the  table  in  Part  I  §  258. 

A.  Nutrients.     Apply  tests  for  carbohydrate,  fat,  and  pro- 
tein to  various  kinds  of  food.      In  testing  for  fat,  shake  the 
crushed  food  with  gasoline,  pour  off  the  gasoline  into  a  dish, 
let  it  evaporate,  and  examine  the  residue.      (Keep  gasoline 
away  from  flames.)      State  the  results  in  each  case. 

B.  Water.    Proceed  as  in  Exp.  14  with  various  kinds  of  food. 

C.  Mineral  Matter.     Heat  a  sample  in  an  evaporating  dish 
or  on  a  piece  of  porcelain  until  all  traces  of  carbon  are  removed. 
The  white  or  whitish  residue  is  mineral  matter.    Further  tests 
may  be  made  for  a  chloride  or  sulphate  or  for  different  metals, 
e.g.  sodium  (flame  test),  calcium   (Exp.  113  H),  aluminium 
(Exp.  209  (c) ),  potassium  (flame  test). 

Experiment  120  —  Testing  Flour 

MATERIALS.  —  Flour,  cheese  cloth. 

A.  Fat.  Apply  a  test  for  fat  to  dry  flour.  State  the 
result. 


io6  CHEMISTRY 

B.  Carbohydrate  and  Protein.    Put  a  little  flour  on  a  small 
piece  of  cheese  cloth  and  tie  the  cloth  into  a  bag.     Put  the 
closed  bag  in  an  evaporating  dish  half  full  of  water  and  move 
the  bag  about  in  the  water,  squeezing  it  occasionally.    Finally, 
remove  the  bag  and  wash  it  thoroughly  in  running  water. 
Apply  tests  for  starch  and  protein  to  (i)  the  solid  left  in 
the  bag  and  (2)  the  solid  that  settles  in  the  dish.    State  each 
result. 

C.  Water  and  Mineral  Matter.    Proceed  as  in  Exp.  119  B, 
and  C.     State  each  result. 

D.  Devise  an  experiment   to  show  that  carbon  dioxide  is 
formed  during  bread  making.      Before  proceeding,  submit  the 
details  to  the  Teacher.      (Suggestion.  —  See  Exp.  127  I.) 

SUPPLEMENTARY  EXPERIMENTS 

Experiment  121  —  Preparation  of  Invert  Sugar  (Dextrose 
and  Levulose)  from  Sucrose 

MATERIALS.  —  Cane  sugar,  Fehling's  solution. 

Add  a  few  drops  of  concentrated  hydrochloric  acid  to  about  25 
cubic  centimeters  of  cane  sugar  solution  and  boil  several  minutes. 
Neutralize  with  sodium  hydroxide  solution  and  test  with  Fehling's 
solution.  State  the  result. 

Experiment  122  —  Testing  for  Sugar  in  Vegetables 
and  Fruits 

MATERIALS.  —  Fehling's  solution  and  the  substances  enumerated  below. 
Apply  Fehling's  test  to  a  clear  solution  obtained  from  each  of  the 
following:  Apple,  banana,  orange,  carrot,  turnip,  raisins,  and  other 
available  vegetables  and  fruits.  Prepare  the  solution  by  cutting  or 
grinding  the  substance  into  very  small  pieces,  or  adding  a  little 
water,  and  squeezing  the  soft  mass  in  a  piece  of  cheese  cloth.  State 
the  result  in  each  case. 

Experiment  123  —  Detection  of  Starch  by  Iodine 

MATERIALS.  —  Dilute  solution  of  iodine,   mortar  and  pestle,  potato, 

rice,  bread. 

(a)  Test  potato,  rice,  and  bread  for  starch  by  grinding  a  little 
of  each  separately  with  water  in  a  mortar,  and  then  adding  a  few 


ORGANIC   COMPOUNDS  —  FOOD  107 

drops  of  the  extract  to  a  very  dilute  solution  of  iodine.    State  the  result 
in  each  case. 

(b)  Proceed  with  the  testing  as  in  (a),  using  substances  not  posi- 
tively known  to  contain  starch,  such  as  baking  powder,  leaves  of 
different  kinds  of   trees,  roots  of  vegetables,  popped   corn,  straw. 
State  the  result. 

Experiment  124  —  Conversion  of  Starch  into  Sugar 
by  an  Enzyme 

MATERIALS.  —  Cracker  or  bread,  Fehling's  solution. 

(a)  Grind  a  small  piece  of  cracker  or  bread  in  a  mortar  with  a 
little  water,  and  test  the  mixture  with  Fehling's  solution.  State  the 
result. 

(6)  Chew  a  piece  of  cracker  or  bread  for  a  minute  or  two,  add  a 
little  water,  and  test  the  clear  solution  with  Fehling's  solution.  Com- 
pare with  the  result  in  (a). 

Experiment  125  —  Properties  of  Dextrin 

MATERIALS.  —  Dextrin,  tannin,  Fehling's  solution,  toasted  bread. 

(a)  Examine  dextrin  and  state  its  characteristic  properties.  Taste 
a  little  and  compare  its  sweetness  with  that  of  sucrose  and  dextrose. 

(6)  Dissolve  dextrin  in  a  little  water  and  note  the  properties  of 
the  solution.  Apply  some  of  the  solution  to  one  side  of  a  piece  of 
paper,  fold  over  the  coated  side,  press  together,  and  examine  after  a 
short  time.  State  the  final  result.  Dilute  the  rest  of  the  solution 
and  use  in  (c)  and  (d). 

(c)  Add  tannin  to  part  of  the  dextrin  solution  from  (b).    Observe 
and  state  the  result. 

(d)  Apply  Fehling's  test  to  dextrin  solution.     State  the  result. 

(e)  Soak  toasted  bread  in  water,  filter,  and  test  the  filtrate  for 
dextrin  (as  in  (c)  and  (d) ). 

Experiment  126  —  Properties  of  Guncotton,  Collodion, 
and  Celluloid 

MATERIALS.  —  Guncotton,  cotton,  collodion,  celluloid. 

(a)  Examine  guncotton  and  state  its  obvious  properties.  Rub 
it  between  the  fingers  and  compare  the  feeling  with  that  produced 
by  ordinary  cotton.  Burn  a  little  guncotton  and  ordinary  cotton,  and 
compare  the  results. 


io8  CHEMISTRY 

(6)  Pour  or  brush  a  little  collodion  solution  on  a  glass  plate.  As 
soon  as  the  solvent  has  evaporated,  examine  and  describe  the  result- 
ing film  of  collodion.  Ignite  a  small  piece  of  the  film,  and  observe 
and  state  the  result. 

(c)  Examine  celluloid  and  state  its  characteristic  properties,  espe- 
cially the  odor.  Set  fire  to  a  small  piece,  and  observe  and  state  the 
result. 

Experiment  127  —  Preparation  and  Properties  of  Ethyl 
Alcohol 

MATERIALS.  —  Grape  sugar,  yeast,  calcium  hydroxide  solution,  animal 
charcoal,  sodium  hydroxide.  The  apparatus  consists  of  a  large 
bottle  provided  with  a  one-hole  stopper  fitted  with  a  delivery  tube 
which  reaches  to  the  bottom  of  a  small  bottle. 

I.  Preparation.     Put  500  cubic  centimeters  of  water  in  the  bottle, 
add  60  grams  of  grape  sugar,  and  shake  until  dissolved.    Break  a  fresh 
yeast  cake  into  small  pieces,  grind  it  to  a  paste  with  a  little  water, 
and  add  it  to  the  sugar  solution.    Fill  the  small  bottle  half  fulj  of 
calcium  hydroxide  solution,  and  carefully  cover  this  solution  with  a 
little  kerosene.    Stand  the  apparatus  where  the  temperature  is  25° 
to  30°  C. 

Fermentation  begins  at  once,  and  carbon  dioxide  —  one  of  the 
products  —  bubbles  through  the  calcium  hydroxide  solution,  which  is 
protected  from  the  action  of  carbon  dioxide  in  the  air  by  the  kerosene. 
The  operation  should  be  allowed  to  continue  three  or  four  days. 
The  alcohol  must  be  separated  by  distillation. 

II.  Properties.     The  distillation  is  performed  with  the  apparatus 
used  in  Exp.  26.    Fill  the  flask  half  full  of  the  liquid  from  I,  add  a  few 
pieces  of  pipestem  (or  granulated  zinc,  or  glass  tubing)  to  prevent 
"bumping,"  and  distil  about  50  cubic  centimeters.     Save  the  dis- 
tillate.    Replace  the  residue  in  the  flask  by  more  liquid   from  I, 
distil  again,  and  repeat  this  operation  until  all  the  liquid  has  been 
used.    Replace  the  one-hole  stopper  with  a  two-hole  stopper,  insert 
the  bent  tube  into  one  hole  and  a  thermometer  into  the  other  so  that 
the  bulb  just  touches  the  surface  of  the  combined  distillates,  which 
should  now  be  distilled.     Heat  gently,  and  collect  in  a  separate 
receiver  the  distillate  which  is  formed  when  the  temperature  reaches 
about  95°  C.    This  distillate  contains  most  of  the  alcohol.     Test  as 
follows:   (a)  Note  the  odor. 


ORGANIC   COMPOUNDS  —  FOOD  109 

(b)  Drop  a  little  into  a  warm  dish,  and  hold  a  lighted  match  over 
it.    If  it  does  not  burn,  it  shows  that  the  alcohol  is  too  dilute.    Put 
a  little  in  a  dish,  warm  gently,  and  light  the  vapor.     Describe  the 
result. 

(c]  To  the  remainder  add  a  crystal  or  two  of  iodine  and  just  enough 
sodium  hydroxide  solution  to  dissolve  the  iodine.     Warm  gently 
several  minutes  and  then  cool.    The  yellow  product  is  iodoform  and 
its  formation  is  a  test  for  alcohol. 

Experiment  128  —  Preparation  and  Properties  of 
Formaldehyde 

MATERIALS.  —  Methyl   alcohol,    copper  wire,  forceps,  formalin,  silver 
nitrate  solution. 

(a)  Put  a  few  cubic  centimeters  of  methyl  alcohol  in  a  test  tube  and 
stand  the  test  tube  in  a  rack.    Wind  a  piece  of  copper  wire  into  a  spiral 
around  a  glass  rod  or  lead  pencil.    Slip  the  spiral  from  the  rod,  grasp 
one  end  with  the  forceps,  and  heat  the  wire  red-hot  in  the  flame. 
Then  quickly  drop  it  into  the  methyl  alcohol.     The  pungent  vapor 
which  is  suddenly  produced  is  largely  the  vapor  of  formaldehyde. 

(b)  Proceed  as  in  Exp.  105  (/),  using  formalin  instead  of  glu- 
cose solution. 

(c}  Let  gelatin  or  albumin  stand  in  formalin  for  several  days. 
Observe  and  state  the  change  in  the  solid. 

Experiment  129  —  Properties  of  Ether 

MATERIALS.  —  Ether,  evaporating  dish,  glass  plate,  wax. 

Precaution.     Ether  vapor  is  easily  ignited,  and  ether  should  never  be 
brought  near  a  flame  unless  the  directions  so  state. 

(a)  Pour  a  little  ether  into  a  dish  or  test  tube  and  observe  the  odor 
and  volatility.    Taste  cautiously.    Pour  a  drop  upon  a  glass  plate  or 
a  block  of  wood.     How  does  its  rate  of  evaporation  compare  with 
that  of  alcohol?    Pour  a  little  upon  the  hand  and  describe  the  result. 

(b)  Add  a  bit  of  wax  to  a  few  cubic  centimeters  of  ether,  and  shake 
well.    Observe  the  result.    If  the  result  is  doubtful,  pour  the  liquid 
carefully  upon  a  glass  plate,  and  observe  the  final  result.     Draw  a 
conclusion  regarding  the  solvent  power  of  ether. 

(c)  Put  a  few  drops  of  ether  in  an  evaporating  dish,  and  cautiously 
bring  a  Bunsen  flame  near  it.    Describe  the  result. 


SULPHUR -SULPHUR   COMPOUNDS 

Experiment  130  —  Physical  Properties  of  Sulphur 

(//  desired,  (c)  may  be  omitted  and  performed  later  as  Exp.  132.) 
MATERIALS.  —  Sulphur,  graduated  cylinder. 

(a)  Examine  specimens  of  brimstone  and  flowers  of  sulphur, 
and  state  the  characteristic  properties  of  each. 

(b)  Determine  the  specific  gravity  of  sulphur  by  the  method 
given  in  Exp.  88  (b*). 

(c)  Fill  a  test  tube  one  fourth  full  of  small  lumps  of  sulphur 
and  heat  carefully  until  all  the  sulphur  is  melted.     Observe 
the  color  and   consistency  of  the  melted  sulphur.     Increase 
the  heat,  and  observe  as  before.     Continue  to  heat  until  the 
sulphur  boils  and  then  observe  as  before.     Let  the  test  tube 
cool,  and  save  it  for  Exp.  132.     Summarize  the  observations 
made  when  sulphur  was  heated. 

NOTE.  —  If  the  test  tube  should  break  during  the  heating,  extinguish 
the  burning  sulphur  with  sand. 

Experiment  131  —  Preparation  of  Crystallized  Sulphur 

MATERIALS.  —  Sulphur  (roll),  carbon  disulphide,  evaporating  dish. 

A.  Monoclinic.  Fix  a  folded  filter  paper  firmly  in  a  funnel, 
and  place  the  funnel  in  a  test  tube  which  stands  in  a  rack. 
Fill  another  test  tube  two-thirds  full  of  roll  sulphur,  heat  it 
at  first  throughout  its  length,  next  gradually  increase  the  heat 
until  all  the  sulphur  is  melted,  and  then  quickly  pour  it  upon 
the  filter  paper.  Let  it  cool  until  crystals  appear  just  below 
the  surface,  and  then  pour  out  the  remaining  melted  sulphur. 
Remove  the  paper  and  adhering  sulphur,  and  cut,  or  break, 
open  the  cone  of  crystallized  sulphur.  Observe  and  record 
the  properties  of  the  crystals,  especially  the  shape,  size,  color, 
luster,  brittleness,  and  any  other  characteristic  property. 


SULPHUR  —  SULPHUR   COMPOUNDS         in 

Allow  the  best  crystals  to  remain  undisturbed  for  a  day  or 
two;   then  examine  again,  and  record  any  marked  changes. 

B.  Orthorhombic.  Put  about  3  grams  of  powdered  roll 
sulphur  in  a  test  tube  and  add  about  10  cubic  centimeters  of 
carbon  disulphide  —  remember  to  keep  the  carbon  disulphide 
away  from  flames.  Shake  until  most  of  the  sulphur  is  dis- 
solved, then  filter  the  solution  (or  pour  the  clear  liquid)  into 
an  evaporating  dish  to  crystallize.  It  is  advisable,  and  often 
absolutely  necessary,  to  stand  the  dish  in  the  hood  or  out  of 
doors,  where  there  is  no  flame  and  where  the  offensive  vapor 
will  be  quickly  removed.  Allow  the  liquid  to  evaporate; 
watch  the  crystallization  toward  the  end,  if  convenient,  and 
when  the  liquid  has  evaporated  almost  entirely,  remove  and 
dry  the  best  crystals.  Examine  them  as  in  A  and  record 
their  properties. 

Experiment  132  —  Preparation  of  Amorphous  Sulphur 

MATERIALS.  —  Sulphur,  test  tube,  evaporating  dish. 

Put  a  few  pieces  of  roll  sulphur  in  a  test  tube  (see  Exp. 
132  (c)),  heat  carefully  until  the  sulphur  boils,  and  then  quickly 
pour  the  molten  sulphur  into  a  dish  of  cold  water.  This  is  the 
plastic  variety  of  amorphous  sulphur.  Note  its  properties. 
Preserve,  and  examine  it  after  twenty-four  hours.  Describe 
it,  and  compare  its  properties  with  those  previously  observed. 
Pulverize  a  small  piece  and  test  its  solubility  in  carbon  disul- 
phide. State  the  result. 

Experiment  133  —  Chemical  Properties  of  Sulphur 

MATERIALS.  —  Sulphur,  deflagrating  spoon,  bottle,  iron  thread. 

(a)  Set  fire  to  a  little  sulphur  in  a  deflagrating  spoon,  and 
lower  the  spoon  into  a  bottle.     Cautiously  waft  the  fumes 
toward  the  nose,  and  observe  and  describe  the  odor.     What 
is  the  product  of  burning  sulphur?    What  does  its  formation 
show  about  the  combining  power  of  sulphur? 

(b)  Fill  a  test  tube   one-fourth   full  of  sulphur  and  press 
iron  thread  down  upon  the  sulphur  until  the  test  tube  is  nearly 


112 


CHEMISTRY 


full.  Heat  the  test  tube  strongly  until  the  sulphur  boils  or 
there  is  marked  evidence  of  chemical  action.  Remove  the 
test  tube  from  the  flame  as  soon  as  the  reaction  begins.  Ob- 
serve and  describe  the  result.  What  is  the  name  of  the  product 
of  the  chemical  change? 

(c)  Expose  moist  sulphur  to  the  air  for  an  hour  or  more. 
Shake  with  water,  filter,  and  apply  the  test  for  a  sulphate  to 
the  filtrate.  State  the  result.  What  property  of  sulphur  is 
illustrated  by  this  experiment? 

Experiment  134  —  Preparation  and  Properties  of  Sulphur 
Dioxide  and  Sulphurous  Acid 

MATERIALS.  —  Sodium  sulphite,  concentrated  sulphuric  acid,  litmus 
paper,  three  bottles,  two  glass  plates,  joss  stick,  colored  flower. 
The  apparatus  is  shown  in  Fig.  140. 

I.  Preparation,  (a)  Sulphur  Dioxide.  Put  about  10  grams 
of  sodium  sulphite  in  the  flask,  cover  it  with  water,  and  insert 

the  stopper  with  its  tubes.  Adjust 
the  apparatus  as,  shown  in  Fig.  140. 
Fill  the  cup  with  concentrated  sul- 
phuric acid,  press  the  pinchcock  a 
little,  and  let  the  acid  flow  drop  by 
drop  upon  the  sodium  sulphite. 
Sulphur  dioxide  gas  is  evolved  and 
passes  into  the  bottle,  which  should 
be  removed  when  full,  as  previ- 
ously described.  Moist  blue  litmus 
paper  held  for  an  instant  at  the 
mouth  of  the  bottle  will  show 
(by  change  in  color)  when  the 
latter  is  full.  Collect  two  bottles 
of  gas,  cover  each  with  a  glass 
plate,  and  set  aside  until  needed. 
(b)  Sulphurous  Acid.  As  soon  as  the  second  bottle  of  gas 
has  been  removed  and  covered,  put  in  its  place  a  bottle  one 
fourth  full  of  water.  Adjust  its  height  (if  necessary)  by  wooden 


Fig.  140.  —  Apparatus  for  Pre- 
paring Sulphur  Dioxide. 


SULPHUR  —  SULPHUR   COMPOUNDS          113 

blocks,  so  that  the  end  of  the  delivery  tube  is  just  above  the 
surface  of  the  water.  Continue  to  add  the  acid  drop  by  drop, 
at  intervals.  Shake  the  bottle  occasionally.  Meanwhile 
proceed  as  in  II  with  the  sulphur  dioxide  gas  already  collected. 

II.  Properties  of  Sulphur  Dioxide  Gas.     (a)  Observe  and 
state  the  most  obvious  physical  properties,  e.g.  color,  odor 
(cautiously),  density. 

(b)  Hold  a  blazing  joss  stick  in  the  same  bottle  of  the  gas 
for  a  few  seconds.    Does  the  gas  burn  or  support  combustion? 

(c)  Stand  the  bottle  (used  in  (a)  )  mouth  downward  -in  a 
vessel  of  water.     Shake,  still  keeping  the  mouth  submerged. 
State  the  result.     Test  the  solution  with  litmus  paper.     Is 
the  resulting  liquid  acid,  alkaline,  or  neutral? 

(d)  Moisten  a  colored  flower  with  a  few  drops  of  water, 
hang  it  in  the  remaining  bottle  of  sulphur  dioxide,  holding  it 
in  place  by  putting  the  stem  between  the  glass  and  a  cork. 
Observe  and  describe  any  change  in  the  color  of  the  flower. 
(If  a  flower  is  not  available,  use  colored  paper.) 

III.  Properties  of  Sulphurous  Acid.     Test   as   follows   the 
solution  of  sulphurous  acid  prepared  in  I  (b) :   (a)  Observe  the 
odor  and  the  taste  cautiously.    State  the  result  in  each  case. 

(b)  Apply  the  litmus  test,  and  state  the  result. 

(c)  Divide  the  solution  into  two  parts.     Save  one  for  (d). 
Into  the  other  put  a  piece  of  magnesium.    State  the  result. 

(d)  Pour  a  few  drops  of  potassium  permanganate  solution 
into  the  other  portion  of  the  sulphurous  acid  solution.     Ob- 
serve and  state  the  result.     What  chemical  change  has  the 
sulphurous  acid  undergone?    If  in  doubt,  suggest  an  experi- 
ment which  will  answer  the  question. 

Experiment  135  —  Properties  of  Sulphuric  Acid 

MATERIALS.  —  Concentrated  sulphuric  acid,  small  graduated  cylinder, 
hydrometer,  thin  stick  of  wood,  sugar. 

(a)  Weigh  a  25  cubic  centimeter  graduated  cylinder,  pour 
in  concentrated  sulphuric  acid  to  a  convenient  height,  and 
weigh  again.  Read  the  volume  of  the  acid.  From  the  weight 


ii4  CHEMISTRY 

and  volume  of  the  acid,  calculate  its  specific  gravity.  Verify 
the  result  by  reading  the  hydrometer  which  floats  in  a  sample 
of  the  same  acid.  (This  apparatus  should  be  arranged  for  the 
class  by  the  Teacher.) 

(b)  Add  an  equal  volume  of  concentrated  sulphuric  acid  to 
a  test  tube  one-fourth  full  of  water,  and  observe  the  change 
in  temperature.    Save  the  solution  for  (c)  and  (d). 

(c)  (i)  Write  some  letters   or  figures   with  the   sulphuric 
acid  from  (b)  on  a  sheet  of  white  paper,  and  move  the  paper 
back'  and  forth  over  a  low  flame,  taking  care  not  to  set  fire  to 
the  paper.    As  the  water  evaporates  the  dilute  acid  becomes 
concentrated.     Observe  and  describe   the  result.     (Paper  is 
largely  a  compound  of  carbon,  hydrogen,  and  oxygen,  and  the 
hydrogen  and  oxygen  are  present  in  the  proportion  to  form 
water.)     Explain  the  general  chemical  change  in  this  experi- 
ment.    (2)  Warm  the  acid  in  the  test  tube  saved  from  (b), 
stand  a  stick  of  wood  in  the  acid,  and  allow  it  to  remain  for 
fifteen  minutes  or  more.    Then  remove  the  stick  and  wash  off 
the  acid.    Describe  and  explain  the  change  in  the  wood.     (3) 
Proceed  as  in  Exp.  104  (d). 

(d)  Perform  in  hood.    Put  a  few  drops  of  concentrated  sul- 
phuric acid  in  an  evaporating  dish,  support  the  dish  on  a 
gauze-covered  ring  attached  to  an  iron  stand,  and  heat  in- 
tensely.    Observe  and  describe  the  result.     Stop  heating  as 
soon  as  the  result  is  obtained  and  let  the  dish  cool  before 
removing  it. 

Experiment  136  —  Tests  for  Sulphuric  Acid,  Sulphates, 
and  SO4-ions 

MATERIALS.  — •  Sulphuric  acid,  sodium  sulphate,  barium  chloride  solu- 
tion, calcium  sulphate,  charcoal,  powdered  charcoal,  blowpipe, 
silver  coin. 

A.  Sulphuric    Acid.     Recall   a   test   for   concentrated   sul- 
phuric acid.    How  could  the  same  test  be  utilized  in  the  case 
of  dilute  sulphuric  acid? 

B.  Sulphuric    Acid    and    Soluble    Sulphates,   i.e.   solutions 


SULPHUR  —  SULPHUR   COMPOUNDS          115 

containing  SO4-ions.  Add  barium  chloride  solution  to  the 
solution  of  the  acid  or  the  sulphate,  and  boil  with  dilute 
hydrochloric  acid.  If  no  sulphur  dioxide  gas  is  liberated  and 
an  insoluble  precipitate  remains,  then  the  original  solution 
contained  SO4-ions.  (See  Exp.  134.) 

C.   Insoluble   Sulphates.     Proceed   as   in   Exp.    97   A    (#), 
using  calcium  sulphate  (or  any  insoluble  sulphate). 

SUPPLEMENTARY  EXPERIMENTS 

Experiment  137  —  Sulphur  Matches 

(a)  Examine  a  sulphur  match.     Do  you  detect  any  sulphur? 
Where? 

(b)  Light  a  sulphur  match,  and  observe  the  entire  action,  as  far 
as  the  sulphur  is  concerned.     Describe  it. 

(c)  What  is  the  function  of  the  sulphur  in  a  burning  match? 

Experiment  138  —  Preparation  and  Properties  of 
Hydrogen  Sulphide 

MATERIALS.  —  Ferrous  sulphide,  dilute  hydrochloric  acid,  three  bottles, 
three  glass  plates,  stoppered  bottle,  litmus  paper.     The  apparatus 
is  shown  on  Fig.  141. 
Precaution.     Hydrogen  sulphide  is  a  poisonous  gas  and  has  an  offensive 

odor.     It  should  not  be  inhaled  nor  allowed  to  escape  into  the  laboratory. 

Perform  in  the  hood  all  experiments  with  hydrogen  sulphide. 

I.  Preparation.     Construct  and  arrange  an  apparatus  like  that 
shown  in  Fig.  141.    Fill  the  bottle  A  one  fifth-full  of  coarsely  pow- 
dered ferrous  sulphide,  insert  the  stopper  tightly,  and  adjust  the 
apparatus  so  that  the  end  of  the  delivery  tube  will  be  under  the  sup- 
port of  the  pneumatic  trough.    Introduce  a  little  dilute  hydrochloric 
acid  through  the  dropping  tube.    Hydrogen  sulphide  gas  is  rapidly 
evolved.    If  the  evolution  of  gas  slackens  or  stops,  add  more  hydro- 
chloric acid.     Collect  three  bottles,  removing  each  as  soon  as  full 
and  covering  with  a  glass  plate.    Set  aside  until  needed.    When  all 
the  bottles  have  been  filled  with  gas,  proceed  at  once  with  II. 

II.  Properties.     Study  as  follows  the  hydrogen  sulphide  gas  pre- 
pared in  I:    (a)  Waft  a  very  little  of  the  gas  cautiously  toward  the 
nose,  and  describe  the  odor.     This  odor  is  characteristic  of  hydro- 
gen sulphide,  and  is  a  decisive  test.    Has  the  gas  color? 


n6 


CHEMISTRY 


(b)  Test  the  gas  from  the  same  bottle  with  both  kinds  of  moist 
litmus  paper.    Is  hydrogen  sulphide  acid,  alkaline,  or  neutral? 

(c)  Hold  a  lighted  match  to  the  mouth  of  the  same  bottle.    Ob- 
serve the  color  of  the  flame.    Observe  cautiously  the  odor  of  the  prod- 


Fig.  141.  —  Apparatus  for  Preparing  Hydrogen  Sulphide. 

uct  of  the  burned  gas;   to  what  compound  is  the  odor  due?    What, 
then,  is  one  constituent  of  hydrogen  sulphide? 

(d)  Burn  another  bottle  of  hydrogen  sulphide  and  hold  a  cold 
bottle  over  the  burning  gas.    What  additional  experimental  evidence 
does  this  result  give  regarding  the  composition  of  hydrogen  sulphide? 

(e)  Repeat  any  of  the  above  with  the  remaining  bottle  of  the  gas. 
Required  Exercises.  —  i.  Summarize  briefly  the  properties  of  hydro- 
gen sulphide  gas. 

2.  State  the  experimental  evidence  of  its  composition. 

Experiment  139  —  Preparation  and  Properties  of  Sulphides 

MATERIALS.  —  Hydrogen  sulphide  water,  clean  copper  wire,  clean 
sheet  lead,  bright  silver  coin,  lead  oxide  (litharge);  solutions  of 
lead  nitrate,  arsenic  trioxide  (in  hydrochloric  acid),  tartar  emetic, 
zinc  sulphate. 

(a)  Obtain  a  bottle  half  full  of  hydrogen  sulphide  water,  and  hold 
successively  at  the  mouth,  or  in  the  neck,  of  the  bottle  (i)  a  clean 
copper  wire,  (2)  a  bright  strip  of  lead,  and  (3)  an  untarnished  silver 


SULPHUR  — SULPHUR    COMPOUNDS          117 

coin.    Describe  the  result  in  each  case.    These  compounds  are  sul- 
phides of  the  respective  metals;  give  the  name  of  each. 

(b)  Put  a  little  litharge  —  the  brownish  yellow  oxide  of  lead  —  in 
a  test  tube,  cover  it  with  hydrogen  sulphide  water,  and  warm  gently. 
The  product  is  lead  sulphide.     Describe  it.     Explain  the  chemical 
change. 

(c)  Add  hydrogen  sulphide  water  to  lead  nitrate  solution.     The 
product  is  lead  sulphide.    Observe  the  color. 

(d)  Proceed  as  in  (c)  with  the  arsenic  solution.    Observe  the  color 
of  the  arsenic  sulphide. 

(e)  Proceed  as  in  (c)  with  the  tartar  emetic  solution.     Tartar 
emetic  is  a  compound  of  antimony.    Observe  the  color  of  the  anti- 
mony sulphide. 

(/)  Proceed  as  in  (c)  with  the  zinc  sulphate  solution.     Observe 
the  color  of  the  zinc  sulphide. 

Experiment  140  —  Properties  of  Sulphurous  Acid 

Prepare  a  solution  of  sulphurous  acid,  or  obtain  some  from  the 
Teacher,  and  proceed  as  follows:  — 

(a)  Put  about   15   cubic  centimeters  of  sulphurous  acid  into  an 
evaporating  dish,  support  the  dish  on  a  gauze-covered  ring  attached 
to  an  iron  stand  in  the  hood,  heat  gradually  and  note  the  odor  of  the 
liberated  gas.    Blow  the  gas  out  of  the  dish  frequently,  and  then  smell 
of  the  liquid.    Boil  until  most  of  the  liquid  is  evaporated,  and  test 
the  remainder  with  litmus  paper.    What  is  the  effect  of  heat  upon  the 
sulphurous  acid? 

(b)  Put  15  cubic  centimeters  of  sulphurous  acid  into  a  bottle,  and 
let  it  stand  exposed  to  the  air  for  several  days.    Add  a  little  water, 
boil  a  minute  or  two,  and  then  test  the  solution  for  a  sulphate. 

Experiment  141  —  Tests  for  Sulphur 

MATERIALS.  —  Sulphur,  iron  sulphide  (ferrous  sulphide),  a  soluble  sul- 
phate, calcium  sulphate,  albumin. 

A.  Free  Sulphur.    Burn  a  little  sulphur  in  a  deflagrating  spoon 
or  on  the  end  of  a  glass  rod.    Observe  the  color  of  the  flame  and  the 
odor  of  the  gaseous  product. 

B.  In  Sulphides.     See  Exps.  138,  139. 

C.  In  Sulphates.     See  Exp.  136. 

D.  In  Organic  Compounds.     See  117  C,  D. 


BORAX  — BORIC   ACID 

Experiment  142  —  Preparation  of  Crystallized  Borax 
MATERIALS.  —  Borax,  thread. 

Prepare  about  50  cubic  centimeters  of  a  hot,  concentrated 
solution  of  borax.  Pour  the  clear  liquid  into  an  evaporating 
dish,  and  let  the  solution  cool.  Crystals  of  borax  will  form; 
well-shaped  crystals  may  be  obtained  by  suspending  a  piece 
of  thread  in  the  solution  and  removing  it  with  the  adhering 
crystals  before  the  water  entirely  evaporates.  Remove  and 
dry  the  crystals. 

Experiment  143  —  Properties  of  Borax 

MATERIALS.  —  Borax;  crystallized  borax  for  (6). 

(a)  Dissolve  a  little  borax  in  water,  drop  a  piece  of  red 
litmus  paper  into  the  solution,  and  let  the  whole  stand  about 
ten  minutes.    Observe  and  explain  the  result. 

(b)  Test  borax  crystals  and  borax  powder  for  water  of  crys- 
tallization; and  state  the  result.    Expose  to  the  air  for  an  hour 
or  more  some  of  the  borax  crystals  prepared  in  Exp.  142. 

(c)  Apply  the  flame  test  to  a  little  borax  on  the  end  of  a 
clean  test  wire.    What  element  is  contained  in  borax  according 
to  this  test? 

(d)  Dissolve  a  little  borax  in  water,  add  a  few  cubic  centi- 
meters of  ethyl  alcohol  and  of  concentrated  sulphuric  acid, 
and  mix  well.     Test  for  boron  by  dipping  a  clean  test  wire 
into  the  solution  and  holding  it  in  the  outer  part  of  the  Bunsen 
flame.     State  the  result. 

Experiment  144  —  Tests  with  Borax  Beads 

(Each  pupil  need  not  perform  all  of  this  Experiment.) 

MATERIALS.  —  Powdered    borax,    test    wire,    cobalt    nitrate    solution, 
copper  sulphate  solution,  manganese  sulphate  solution. 

Heat  the  looped  end  of  the  clean  test  wire  and  dip  it  into 
powdered  borax.  Heat  the  adhering  borax  in  the  flame, 


BORAX  — BORIC   ACID 


119 


rotating  the  wire  slowly,  until  no  further  change  occurs;  con- 
tinue to  dip  it  into  the  borax  and  heat  in  the  flame  until  a 
small,  more  or  less  transparent,  bead  is  formed. 

A.  Cobalt  Compounds.  Moisten  a  borax  bead  with  cobalt 
nitrate  solution.  Heat  the  bead  in  the  oxidizing  part  of  the 
Bunsen  flame  (Fig.  142);  rotate  the  bead  while  heating  it, 
otherwise  it  may  drop  off  the  wire.  Observe  the  color 
of  the  cold  bead.  If  it  is  black,  melt  a  little  more  borax 
into  the  bead;  if  faintly  colored,  moisten  again  with  the  cobalt 
solution.  The  color  is  readily  detected  by  looking  at  the 
bead  against  a  white  object  in  a  strong  light,  or  by  examining 
it  with  a  lens.  When  the  color  has  been  definitely  determined, 
heat  the  bead  in  the  reducing  flame  (Fig.  142).  Compare 


Fig.  142.  —  Heating  a  Borax  Bead  in  the  Oxidizing  Flame  (left)  and 
the  Reducing  Flame  (right). 

the  color  of  the  cold  bead  with  the  previous  observation. 
State  the  result.  Remove  the  bead  from  the  wire  by  dipping 
it,  white  hot,  into  water;  the  sudden  cooling  shatters  the  bead, 
which  may  then  be  easily  rubbed  or  scraped  from  the  wire. 

B.  Copper  Compounds.     Make   another  bead  on   a   clean 
wire,  moisten  it  with  copper  sulphate  solution  and  heat  it  in 
the  oxidizing  flame;  and  then  in  the  reducing  flame.    Com- 
pare the  colors  of  the  cold  beads,  and  state  the  result. 

C.  Manganese  Compounds.     Make  another  bead  on  a  clean 
wire,  moisten  it  with  manganese  sulphate  solution,  and  pro- 
ceed as  in  B.    Compare  the  colors  of  the  cold  beads,  and  state 
the  result. 

D.  Miscellaneous.      Obtain    unknown    solutions    from    the 
Teacher  and  test  them  with  a  borax  bead. 


120  CHEMISTRY 

SUPPLEMENTARY  EXPERIMENTS 

Experiment  145  —  Preparation  and  Properties  of  Boric 

Acid 

MATERIAL.  —  Powdered  borax. 

I.  Preparation.    Heat  about  25  cubic  centimeters  of  water  nearly 
to  boiling  in  a  large  test  tube,  and  slowly  add  about  10  grams  of 
powdered  borax;    heat  until  the  borax  is  dissolved.     Ppur  about  5 
cubic  centimeters  of  concentrated  hydrochloric  acid  slowly  into  the 
hot  solution  of  borax,  mix  well  by  stirring,  and  then  stand  the  test 
tube  in  the  test-tube  rack  to  cool,  or  cool  it  in  a  stream  of  water. 
Crystals  of  boric  acid  will  separate  from  the  solution.     Filter  (with 
a  filter  pump,  if  one  is  available),  and  wash  the  crystals  while  upon 
the  filter  paper  with  a  little  cold  water.    Redissolve  a  portion  of  the 
crystals  in  a  very  small  quantity  of  boiling  water,  and  let  the  solu- 
tion cool  slowly.     Later  examine  the  crystals  for  crystal  form  and 
luster. 

II.  Properties,     (a)  Examine  a  specimen,  and  state  the  proper- 
ties, e.g.  crystal  form,  color,  luster,  and  the  feeling  when  rubbed 
between  the  fingers. 

(b)  Dissolve  a  little  in  water,  test  the  solution  with  litmus  paper, 
and  state  the  result. 

(c)  Proceed  as  in  Exp.  143  (d),  using  boric  acid  instead  of  borax. 


SILICON  — GLASS 

Experiment  146  —  Properties  of  Silicon 

MATERIALS.  —  Silicon,  sodium  hydroxide,  500  cc.  (or  250  cc.)  graduate. 

(a)  Examine  a  specimen  of  silicon.    Observe  and  state  the 
color,    luster,  texture,  brittleness,  hardness    (compared   with 
glass),  and  any  other  characteristic  physical  property. 

(b)  Determine  the  specific  gravity  by  the  method  described 
in  Exp.  88  (b).    State  the  result. 

(c)  Prepare  a  concentrated  solution  of  sodium  hydroxide 
by  dissolving  about  8  grams  of  the  solid  in  10  cubic  centimeters 
of  water.   Add  about  i  gram  of  powdered  silicon,  heat  the  mix- 
ture to  boiling,  and  test  the  escaping  gas  with  a  blazing  joss 
stick  or  lighted  match.    What  is  the  gas? 

Experiment  147  —  Test  for  Silicon 

MATERIALS.  —  Lead  dish,  powdered  calcium  fluoride,  sand,  test  wire. 

Put  a  little  sand  and  calcium  fluoride  in  a  lead  dish,  add  a 
little  concentrated  sulphuric  acid,  and  stir  with  a  match  until 
well  mixed.  Dip  the  looped  end  of  the  test  wire  into  water  so 
as  to  form  a  film  of  water  within  the  loop,  and  hold  the  loop 
at  several  points  near  the  mixture  in  the  dish  until  the  water 
becomes  white.  If  no  change  occurs,  stir  the  mixture  and 
hold  the  loop  over  the  place  where  there  is  evidence  of  chemical 
action.  What  is  the  white  substance  in  the  water? 

State  in  words  the  chemical  changes  that  led  to  the  forma- 
tion of  the  white  substance  in  the  loop.  Write  the  equations 
for  these  changes. 

SUPPLEMENTARY  EXPERIMENTS 

Experiment  148  —  Preparation  and  Properties  of 

Silicic  Acid 

MATERIALS.  —  Sodium  silicate  solution,  hydrochloric  acid. 

Put  10  cubic  centimeters  of  sodium  silicate  solution  in  an  evapo- 
rating dish,  and  add  10  or  15  cubic  centimeters  of  dilute  hydrochloric 


122  CHEMISTRY 

acid,  stirring  constantly.  The  jellylike  precipitate  is  silicic  acid. 
Rub  some  between  the  fingers  and  state  the  result.  Stand  the  dish 
on  a  gauze-covered  ring  attached  to  an  iron  stand  and  evaporate  the 
solution  to  dryness  in  the  hood.  As  the  mass  hardens,  stir  it  with  a 
glass  rod.  Toward  the  end,  add  more  hydrochloric  acid  and  evaporate 
to  complete  dryness.  Then  heat  intensely  for  five  minutes.  The 
residue  is  silicon  dioxide  mixed  with  chlorides  of  sodium  and  potas- 
sium. When  the  dish  is  cool,  add  about  50  cubic  centimeters  of 
water,  stir  well,  and  filter;  wash  the  residue  once  or  twice,  dry  it, 
remove  as  much  as  possible  from  the  paper,  and  heat  it  carefully  in 
an  evaporating  dish  for  about  five  minutes.  This  residue  is  largely 
silicon  dioxide.  Rub  some  between  the  fingers  or  across  a  glass 
plate.  Is  any  grit  detected?  Collect  some  within  the  loop  of  a  test 
wire  and  heat  it  intensely  in  the  flame  for  several  minutes.  State  the 
result. 

State  the  chemical  changes  that  occur  in  changing  sodium  silicate 
into  the  final  residue. 

Experiment  149  —  The  Cycle  of  Silicon  Dioxide 

MATERIALS.  —  Powdered  silicon  dioxide,  sodium  carbonate,  test  wire. 

Grind  about  i  gram  of  silicon  dioxide  to  a  very  fine  powder  (or 
obtain  the  powdered  substance  from  the  Teacher).  Heat  about  8 
grams  of  crystallized  sodium  carbonate  in  an  evaporating  dish  until 
the  water  of  crystallization  has  been  driven  off.  Mix  the  silicon 
dioxide  and  anhydrous  sodium  carbonate  thoroughly  by  grinding 
them  together  in  a  mortar.  Heat  the  looped  end  of  the  test  wire, 
dip  it  into  the  mixture  and  heat  the  substance  in  the  flame,  rotating 
the  wire  slowly  as  in  the  preparation  of  a  borax  bead;  continue  to 
dip  the  bead  into  the  mixture  and  to  heat  intensely  until  a  moderate 
sized  bead  is  formed.  Heat  this  bead  until  there  is  no  further  evidence 
of  chemical  action.  While  the  bead  is  still  soft,  shake  it  from  the 
wire  into  a  mortar.  Prepare  five  or  six  beads  in  the  same  way,  and 
powder  them.  Transfer  the  powder  to  a  test  tube,  add  a  little  water, 
and  boil;  filter,  if  the  solution  is  not  clear.  Add  dilute  hydrochloric 
acid,  drop  by  drop,  shaking  constantly,  until  the  solution  is  strongly 
acid.  Evaporate  this  acid  solution  to  dryness  in  the  hood,  and 
proceed  from  this  point  as  in  Exp.  148. 

State  the  chemical  changes  by  which  the  silicon  dioxide  was  trans- 
formed into  the  final  substance. 


SILICON  —  GLASS  1 23 

Experiment  150  —  Testing  for  Silicon 

MATERIALS  . — •  Lead  dish,  powdered  calcium  fluoride,  test  wire,  in- 
fusorial earth,  pumice  (powder),  scouring  soap,  glass  (small  frag- 
ments), carborundum  (powder). 

Apply  the  test  for  silicon  to  the  substances  enumerated  above,  as 
in  Exp.  147  (omitting  the  sand,  of  course).  State  the  result  in  each 
case. 

Experiment  151  —  Properties  of  Glass 
MATERIALS.  —  Soft,  hard,  and  flint  glass. 

(a)  Examine  specimens  of  soft,  hard,  and  flint  glass,  and  observe 
their  characteristic  properties. 

(&)  Heat  soft  glass  and  hard  glass  separately  in  the  Bunsen  flame, 
and  observe  and  explain  the  result. 

(c)  Devise  an  experiment  to  show  that  flint  glass  contains  lead, 
Before  proceeding,  submit  the  details  to  the  Teacher. 

(d)  Determine  the  specific  gravity  of  glass  by  the  method  given 
in  Exp.  88  (6). 

(e)  Suggest  a  simple  experiment  to  show  that  glass  contains  silicon. 
(/)  Grind  some  soft  glass  (carefully!)  to  a  fine  powder  (or  obtain 

some  powdered  glass  from  the  Teacher),  transfer  it  to  a  bottle,  fill 
the  bottle  one  fourth  full  of  water,  add  a  few  drops  of  phenol-phthalein 
solution,  cork  the  bottle  tightly,  and  shake  well.  Examine  after  a 
few  days.  State  and  explain  the  result. 


FLUORINE  —  BROMINE  —  IODINE 

Experiment  152  —  Preparation  and  Properties  of 
Hydrogen  Fluoride 

MATERIALS.  —  Lead    dish,    glass    plate,    paraffin,    powdered    calcium 

fluoride,  concentrated  sulphuric  acid. 

Precaution.  Do  not  inhale  hydrogen  fluoride.  It  is  a  corrosive  poison. 
An  aqueous  solution  of  the  gas  —  commercial  hydrofluoric  acid  —  burns 
the  flesh  frightfully.  Perform  this  experiment  with  unusual  care. 

Warm  a  glass  plate  about  10  centimeters  (4  inches)  square 
by  dipping  it  into  hot  water  or  by  moving  it  slowly  above  a 
flame.  Coat  one  surface  uniformly  with  a  thin  layer  of  paraffin 
wax.  Scratch  letters,  figures,  or  a  diagram  through  the  wax 
with  a  pin  or  pointed  glass  rod.  The  wax  should  be  removed 
through  to  the  glass,  and  the  lines  should  be  rather  coarse. 

Put  about  5  grams  of  powdered  calcium  fluoride  in  a  lead 
dish  and  add  just  enough  concentrated  sulphuric  acid  to  form 
a  thin  paste.  Stir  the  mixture  with  a  match.  Hold  a  piece  of 
moist  blue  litmus  paper  in  the  escaping  gas  just  above  the 
surface  of  the  mixture;  state  the  result.  Place  the  glass  plate, 
wax  side  down,  upon  the  lead  dish  and  stand  the  whole  appara- 
tus in  the  hood  for  several  hours,  or  until  some  convenient 
time.  Remove  the  plate  and  scrape  off  the  wax  with  a  knife. 
The  last  portions  can  be  removed  by  rubbing  with  a  cloth 
moistened  with  alcohol  or  turpentine.  Do  not  attempt  to 
melt  off  the  wax  over  the  flame.  If  the  experiment  has  been 
properly  performed,  the  plate  will  be  etched  where  the  glass 
was  exposed  to  the  hydrogen  fluoride  gas.  Write  the  equations 
for  the  essential  chemical  changes  in  this  experiment. 

NOTE.  —  The  lead  dish  should  be  cleaned  in  the  hood  by  scraping 
the  contents  carefully  into  a  waste  jar  and  washing  the  whole  dish  with 
water. 


FLUORINE  —  BROMINE  —  IODINE 


125 


Experiment  153  —  Preparation  and  Properties  of  Bromine 

MATERIALS.  —  Potassium    bromide,    manganese    dioxide,    dilute    sul- 
phuric  acid,    bottle  of    water,  test-tube  holder.      The  apparatus 
(Fig.  143)  consists  of  a  large  test  tube  provided  with  a  one-hole 
rubber  stopper  to  which  is  fitted  the  bent  glass  tube;    the  total 
length  of  the  glass  tube  is  about  30  centimeters  (12  inches). 
Precaution.     Bromine  is  a  corrosive  liquid,  which  forms,  at  the  ordinary 
temperature,  a  suffocating  vapor.     All  experiments  in  which  bromine  is 
used  or  bromine  vapor  is  evolved  should  be  performed  in  the  hood. 

Put  about  10  gm.  of  potassium  bromide  in  the  test  tube, 
add  an  equal  weight  of  manganese  dioxide,  and  also  10  cubic 
centimeters  of  dilute  sulphuric  acid.  Insert  the  stopper  and 
its  tube,  attach  the  test-tube  holder,  and  warm  gently. 
Bromine  vapor  soon  appears  in  the  test  tube  and,  if  the 
heat  is  sufficient,  the  vapor  will  escape  from  the 
delivery  tube.  Regulate  the  heating  so  that 
this  vapor  will  condense  and  collect  in  the  lower 
bend  of  the  delivery  tube.  Both  vapor  and 
liquid  are  bromine.  When  no  further  boiling 
produces  bromine  vapor  in  the  test  tube,  trans- 
fer the  bromine  from  the  delivery  tube  into  a 
bottle  half  full  of  water.  This  operation  can 
be  done  easily  by  holding  the  end  of  the  de- 
livery tube  over  the  mouth  of  the  bottle  and 
heating  the  test  tube  slightly;  the  expanding  pjg  I43._ 
gases  will  force  the  liquid  bromine  out  of  the 
bend  into  the  bottle.  Observe  and  record  the 
physical  properties  of  this  bromine,  especially  the 
color,  solubility  in  water,  specific  gravity,  volatil- 
ity, and  physical  state.  Determine  the  odor  cautiously. 
As  soon  as  these  observations  have  been  made,  cork  the 
bottle  tightly  and  shake  it  vigorously.  Observe  the  result, 
and  draw  a  conclusion  about  the  solubility  of  bromine  in 
water.  Save  the  bottle  and  contents  for  Exp.  154. 

NOTE.  —  Wash  the  delivery  tube  free  from  all  traces  of  bromine, 
taking  care  to  get  none  on  the  hands.  Throw  the  contents  of  the  test 
tube  into  a  waste  jar  in  the  hood  and  wash  the  tube. 


Apparatus 
for  Pre- 
paring 
Bromine. 


126  CHEMISTRY 

Experiment  154  —  Preparation  and  Properties  of 
Magnesium  Bromide 

MATERIALS.  —  Bromine  water  (saved  from  Exp.  153  or  obtained  from 
the  Teacher),  magnesium,  chlorine  water. 

Shake  the  corked  bottle  of  bromine  water  until  most  or  all 
of  the  bromine  is  dissolved.  Remove  the  cork  carefully,  add 
a  little  powdered  magnesium,  insert  the  cork,  and  shake 
well.  Let  the  excess  of  magnesium  settle,  and  observe  the 
result.  If  the  change  is  inconspicuous,  add  more  magnesium, 
and  shake.  Pour  the  liquid  into  a  test  tube,  and  observe  the 
appearance;  compare  it  with  the  color  of  the  original  bromine 
water.  Now  add  chlorine  water  drop  by  drop,  shaking  fre- 
quently, until  a  decided  change  in  color  takes  place.  To  what 
is  this  color  due?  State  the  chemical  changes  that  took  place 
upon  the  addition  of  (a)  magnesium  and  (b)  chlorine  water. 
Write  the  equations  for  these  chemical  changes. 

Experiment  155  —  Tests  for  Bromine  in  Bromides 

MATERIALS.  —  Potassium    bromide,    silver    nitrate    solution,    carbon 

disulphide. 

(a)  Add  a  little  concentrated  sulphuric  acid  to  a  little  potas- 
sium bromide  in  a  test  tube;    warm  slightly  if  the  action  is 
not  marked.     Observe  the  result,'  especially  the  color  of  the 
liquid  or  of  the  vapor  just  above  the  liquid.     What  element 
does  it  suggest? 

(b)  To  a  solution  of  a  bromide,  add  a  little  silver  nitrate 
solution,  and  shake.     Observe  the  properties  of  the  precipi- 
tate, especially  the  color  and  texture.     Determine  the  solu- 
bility in  ammonium  hydroxide  by  warming  a  little  of  the 
precipitate  in  ammonium  hydroxide.    State  the  result.    Com- 
pare the  properties  of  silver  bromide  with  those  of  silver 
chloride  (Exp.  36). 

(c)  To  a  solution  of  a  bromide,  add  a  little  chlorine  water 
and  a  few  drops  of  carbon  disulphide,  and  shake.    The  carbon 
disulphide  will  be  colored  yellow  or  brown  by  the  liberated 
bromine. 


FLUORINE  —  BROMINE  —  IODINE  127 

Experiment  156  —  Preparation  and  Properties  of  Iodine 

MATERIALS.  —  Potassium    iodide,    manganese    dioxide,    mortar    and 
pestle,  concentrated  sulphuric  acid,  funnel,  cotton. 

I.  Preparation.     Grind  together  in  a  mortar  about  10  gm. 
of  potassium  iodide  and  about  twice  this  weight  of  manganese 
dioxide.     Put  the  mixture  in  a  test  tube,  moisten  with  water, 
and  add  a   few  cubic  centimeters  of  concentrated  sulphuric 
acid.    Clamp  the  test  tube  vertically  to  an  iron  stand.    Close 
up  the  inner  end  of  the  stem  of  the  funnel  with  a  small  plug 
of  cotton.    Hold  the  funnel  firmly  over  the  mouth  of  the  test 
tube,  and  heat  the  test  tube  gently.      The  vapor  of  the  liber- 
ated iodine  will  fill  the  test  tube,  and  crystals  may  collect  in 
the  upper  part  of  the  test  tube  and  in  the  funnel.      (If  the 
crystals  collect  in  the  test  tube,  force  them  into  the  funnel  by 
heating  the  test  tube  gently  near  the  top.)      Continue  to  heat 
until  enough  iodine  for  several  experiments  collects  in  the 
funnel.    Scrape  the  crystals  into  a  dish. 

II.  Properties.     Study  the  properties  of  the  iodine  as  fol- 
lows:   (a)  Observe  and  record  the  physical  properties,  espe- 
cially the  color  of  the  solid  and  of  the  vapor,  and  the  odor 
(cautiously).     Determine  the  volatility  by  putting  a  crystal 
or  small  piece  in  a  bottle  and  exposing  to  the  sunlight. 

(b)  Heat  a  crystal  in  a  dry  test  tube,  and  invert  the  test 
tube  when  it  is  half  full  of  vapor.    What  does  the  result  show 
about  the  density  of  iodine  vapor? 

(c)  Touch  a  crystal  with  the  finger.    What  color  is  the  stain? 
Will  water  remove  it?    Will  alcohol?    Will  a  solution  of  potas- 
sium iodide?    What  do  these  results  show  about  the  solubility 
of  iodine? 

NOTE.  —  If  crystals  are  left,  use  them  in  the  next  experiment.     Pre- 
serve the  iodine  in  a  stoppered  bottle,  if  not  used  at  once. 


128  CHEMISTRY 

Experiment  157  —  Tests  for  Free  Iodine 

MATERIALS.  —  Iodine,  potassium  iodide,  carbon  disulphide,  starch. 

Precaution.     Carbon   disulphide  is   inflammable.     It   should   not    be 
used  near  flames. 

(a)  Add  a  few  drops  of  carbon  disulphide  to  a  very  dilute 
solution  of  iodine,  which  can  be  prepared  by  dissolving  a  crystal 
of  iodine  in  potassium  iodide  solution.    Shake  well,  and  observe 
the  color  of  the  carbon  disulphide. 

(b)  Grind  a  very  small  lump  of  starch  in  a  mortar  with  a 
little  water,  pour  the  mixture  slowly  into  about  15  cubic  centi- 
meters of  hot  water,  and  stir  the  hot  liquid.    Allow  it  to  cool, 
or  cool  it  by  holding  the  vessel  in  a  stream  of  cold  water.    Add 
a  few  cubic  centimeters  of  the  cold  starch  solution  to  a  test 
tube  nearly  full  of  water,  and  then  add  a  few  drops  of  dilute 
iodine  solution.     Observe  the  result.     (The  starch  should  be 
colored  blue;   if  the  color  is  black,  pour  out  half  of  the  liquid 
and  add  more  water.) 

State  briefly  the  two  tests  for  free  iodine. 

Experiment  158  —  Tests  for  Iodine  in  Iodides 

MATERIALS.  —  Potassium  iodide,  chlorine  water,  starch,  carbon  disul- 
•phide,  silver  nitrate  solution. 

(a)  Add  a  few  drops  of  carbon -disulphide  to  a  very  dilute 
solution  of  potassium  iodide.     Now  add  several  drops  of  chlo- 
rine water,  and  shake  well.     Observe  and  explain  the  result. 

(b)  Add  a  few  cubic  centimeters  of  cold  starch  solution  to 
a  very  dilute  solution  of  potassium  iodide.     Add  a  few  drops  of 
chlorine  water,  and  shake  well.    Observe  and  explain  the  result. 

(c)  To  a  solution  of  an  iodide,  add  a  little  silver  nitrate 
solution,  and  shake.    Observe  the  properties  of  the  precipitate, 
especially  the  color  and  texture.    Test  the  solubility  of  a  little 
of  the  precipitate  in  ammonium  hydroxide,  and  state  the  result. 
Compare  the  properties  of  silver  iodide  with  those  of  silver 
chloride  and  silver  bromide  (see  Exps.  36,  155  (b)). 

(d)  Proceed  as  in   Exp.   155   (a),   using  potassium    iodide 
instead  of  potassium  bromide. 


PHOSPHORUS  —  ARSENIC  —  ANTIMONY  — 
BISMUTH 

Experiment  159  —  Tests  for  Orthophosphoric  Acid 
and  Orthophosphates 

MATERIALS.  — •  Solutions  of  disodium  phosphate,  silver  nitrate,  am- 
monium molybdate,  ammonium  chloride,  magnesium  sulphate,  and 
orthophosphoric  acid,  bone  ash,  fertilizer. 

(a)  Put  a  little  disodium  phosphate  solution  in  a  test  tube 
and  add  a  little  silver  nitrate  solution.    Observe  and  describe 
the  result.    What  is  the  name  of  the  visible  product?    What 
is  its  formula? 

(b)  Put  5  cubic  centimeters  of  disodium  phosphate  solution 
in  a  test  tube  and  add  one  or  two  cubic  centimeters  of  dilute 
nitric  acid;    add  an  equal  volume  of  ammonium  molybdate 
solution.    Observe  and  describe  the  result.     (Warm,  if  no  pre- 
cipitate  appears.)     The  precipitate  is   ammoniunvphospho- 
molybdate  (  (NH^aPO^i^MoOa,  approximately).  .  Apply  this 
test  to  a  dilute  solution  of  orthophosphoric  acid,  and  state 
the  result. 

(c)  To  magnesium  sulphate  solution  add  successively  solu- 
tions   of    ammonium    chloride,    ammonium    hydroxide,    and 
disodium  phosphate.    Observe  and  describe  the  result.    The 
precipitate  is  ammonium  magnesium  phosphate. 

(d)  Dissolve  a  little  bone  ash  in  warm  dilute  nitric  acid, 
filter,  and  apply  the  ammonium  molybdate  test. 

(e)  Proceed  as  in  (d)  with  a  sample  of  fertilizer. 

Experiment  160  —  Tests  for  Metaphosphoric  Acid 
and  Metaphosphates 

MATERIALS.  —  Solutions  of  metaphosphoric  acid,  silver  nitrate,  and 

albumin;   sodium  ammonium  phosphate. 

(a)  To   a   little   metaphosphoric   acid   solution   add   silver 
nitrate  solution  until  a  definite  change  occurs.    Describe  the 


130  CHEMISTRY 

result.    What  is  the  name  of  the  visible  product?    What  is  its 
formula?     Compare  the  color  with  that  observed  in  Exp.  159. 

(b)  Put  a  few  crystals  of  sodium  ammonium  orthophos- 
phate  (microcosmic  salt)  in  a  clean  porcelain  dish,  stand  the 
dish  on  a  gauze-covered  ring,  and  heat  gently.    While  heating, 
hold  a  moistened  piece  of  red  litmus  paper  over  the  dish;  also 
smell   cautiously   of  the   escaping   substance.     What  gas  is 
liberated?     Increase  the  heat  slowly  and  continue  to  heat 
until  no  further  change  seems  to  take  place.     Let  the  dish 
cool,  and  dissolve  the  residue  in  cold  water.     Test  the  solu- 
tion with  silver  nitrate,   and  state   the  result.     Into   what 
compounds    has    the    sodium    ammonium    phosphate    been 
changed? 

(c)  To  a  little  albumin  solution,   add  a  little  metaphos- 
phoric  acid,  and  shake  well.    Observe  and  describe  the  result. 
(This  test  is  not  applicable  to  metaphosphates.) 


Experiment  161  —  Preparation  and  Properties  of 
Arsenic  Trisulphide 

MATERIALS.  —  Hydrogen  sulphide  water,  solutions  of  arsenic  trichloride, 
ammonium  polysulphide,  ammonium  carbonate. 

Add  hydrogen  sulphide  water  to  a  solution  of  an  arsenic 
compound,  such  as  arsenic  trichloride.  The  precipitate  is 
arsenic  trisulphide.  Describe  it.  Filter,  or  let  the  mixture 
stand  until  the  precipitate  settles,  and  then  pour  off  the  liquid. 
Divide  the  precipitate  into  two  parts.  Add  considerable 
ammonium  polysulphide  solution  to  one  part  of  the  precipi- 
tate, and  shake  well.  Observe  and  describe  the  result.  Now 
add  dilute  hydrochloric  acid  carefully  to  acid  reaction.  Observe 
and  describe  the  final  result.  To  the  other  part  of  the  original 
precipitate  add  considerable  ammonium  carbonate  solution, 
and  shake  well.  Observe  and  describe  the  result.  Now  add 
dilute  hydrochloric  acid  solution  slowly  (to  avoid  loss  by 
effervescence)  to  acid  reaction.  Observe  and  describe  the 
final  result.  Summarize  the  properties  of  arsenic  trisulphide. 


PHOSPHORUS  —  ARSENIC  -  -  ANTIMONY      1 3 1 

Experiment  162  —  Properties  of  Antimony  Trichloride 
Pour  a  little  antimony  trichloride  solution  (prepared  as  in 
Exp.  166  B  or  a  similar  one  obtained  from  the  Teacher) 
into  a  large  volume  of  water.  Observe  the  result.  What 
compound  of  antimony  is  formed?  Add  concentrated  hydro- 
chloric acid  drop  by  drop,  shaking  constantly.  Observe  and 
describe  the  result.  What  compound  of  antimony  is  finally 
formed? 

Experiment  163  —  Preparation  of  Antimony  Trisulphide 
Add  hydrogen  sulphide  water  to  the  solution  from  Exp.  166  B 
(or  to  a  similar  solution  obtained  from  the  Teacher).     Ob- 
serve the  result.     Compare  the  color  of  the  precipitate  with 
the  corresponding  arsenic  compound. 

SUPPLEMENTARY  EXPERIMENTS 

Experiment  164  —  Properties  of  Phosphorus  (Optional) 
Precaution.  Phosphorus  is  a  dangerous  substance.  The  yellow 
variety  is  kept  beneath  water,  and  should  be  cut  under  water  and  handled 
only  when  wet.  Dry  yellow  phosphorus  ignites  readily,  and  a  burn 
caused  by  it  heals  very  slowly.  It  is  advisable  to  touch  yellow  phos- 
phorus only  with  wet  fingers;  a  safer  plan  is  to  grasp  it  firmly  with  wet 
forceps  while  it  is  being  cut  or  transferred.  Unusual  care  should  be 
taken  not  to  leave  pieces  of  yellow  phosphorus  in  dishes  or  deflagrating 
spoons  after  the  experiments  have  been  performed.  Ask  the  Teacher  for 
directions  about  the  disposal  of  unused  phosphorus. 

A.  Yellow  Phosphorus.  Fill  a  porcelain  mortar  half  full  of  water, 
and  ask  the  Teacher  to  put  three  small  pieces  of  yellow  phosphorus 
beneath  the  water.  Stand  the  mortar  where  it  will  not  be  upset. 

(a)  Smell  cautiously  of  the  water  in  which  the  phosphorus  has 
been  placed.    If  no  characteristic  odor  is  detected,  proceed  with  the 
other  experiments,  and  observe  the  odor  later.     Describe  it. 

(b)  Wet  the  forceps,  transfer  a  piece  of  the  phosphorus  to  an  evap- 
orating dish  which  has  been  slightly  warmed  by  the  hand  or  a  low 
flame.     Observe  the  result.     Stand  back  as  soon  as  the  phosphorus 
begins  to  burn.    Add  a  little  cold  water  to  the  residue  in  the  dish, 
test  the  solution  with  litmus  paper  (both  colors),  and  state  the  result. 
What  substance  is  in  the  solution? 


132  CHEMISTRY 

(c)  Fill  a  test  tube  half  full  of  water,  transfer  a  piece  of  the  yel- 
low phosphorus  with  wet  forceps  from  the  mortar  to  the  test  tube. 
Warm  the  test  tube  very  gently  and  observe  the  ease  with  which 
phosphorus  melts.    As  soon  as  the  phosphorus  melts,  stand  the  test 
tube  carefully  in  the  test-tube  rack  and  ascertain  the  temperature  of 
the  water  by  a  thermometer.    Record  the  temperature. 

(NOTE.  —  Read  the  Precaution  above.) 

(d)  Have  ready  a  few  crystals  of  iodine  upon  a  piece  of  paper. 
Transfer  the  remaining  piece  of  phosphorus  from  the  mortar  to  an 
evaporating  dish,  dry  it  quickly  by  touching  it  with  the  end  of  a 
piece  of  tightly  rolled  filter  paper,  and  then  slip  the  iodine  upon  the 
dried  phosphorus.    Stand  back  and  observe  the  result. 

(e)  Smell  of  the  head  of  a  phosphorus  tipped  match.     Compare 
the  odor  with  that  observed  in  (a).    Rub  the  head  of  a  phosphorus 
tipped  match  in  a  dark  place,  and  observe  and  describe  the  result. 

B.  Red  Phosphorus.  Obtain  a  little  red  phosphorus  from  the 
Teacher,  (a)  Examine  the  red  phosphorus  and  observe  its  charac- 
teristic properties. 

(b)  Put  a  very  little  in  a  clean  deflagrating  spoon,  and  heat  it  cau- 
tiously in  a  Bunsen  flame  in  the  hood.  Observe  the  result.  Compare 
with  the  result  observed  in  A  (b).  Proceed  with  the  product  as  in 
A  (b),  and  state  the  result. 

Experiment  165  —  Properties  of  Antimony 

MATERIALS.  —  Antimony,  graduate,  blowpipe,  charcoal. 

(a)  Examine  a  specimen  of  antimony  and  state  its  characteristic 
properties,  such  as  the  color,  luster,  crystalline  appearance,  hardness, 
brittleness. 

(b)  Determine  the  specific  gravity  of  antimony  by  the  method 
described  in  Exp.  88  (b) ;  use  a  25  cubic  centimeter  graduate  and  small 
pieces  of  antimony.    State  the  result. 

(c)  Heat  a  small  piece  of  antimony  on  charcoal  with  the  oxidizing 
blowpipe  flame.    Describe  the  result.    What  is  the  white  product? 

Experiment  166  —  Interaction  of  Antimony  and  Acids 
A.   Nitric  Acid.    Boil  a  little  powdered  antimony  with  concen- 
trated nitric  acid  in  the  hood.     Observe  the  effect  on  the  antimony 
What  compound  of  antimony  is  formed? 


ARSENIC  —  ANTIMONY  --  BISMUTH          133 

B.  Aqua  Regia.  Boil  a  little  powdered  antimony  with  aqua  regia 
for  several  minutes  in  the  hood.  Observe  the  result.  What  com- 
pound of  antimony  is  formed?  Pour  off  the  solution  from  any  un- 
changed antimony.  (The  solution  may  be  used  in  Exps.  162,  163.) 

Experiment  167  —  Properties  of  Bismuth 
Proceed  as  in  Exp.  165,  using  bismuth  instead  of  antimony. 

Experiment  168  —  Preparation  and  Properties  of 
Bismuth  Trichloride 

A.  Proceed  as  in  Exp.  166  B,  using  bismuth  instead  of  antimony 
and  taking  care  to  boil  off  most  of  the  acid. 

B.  Proceed  with  the  solution  from  A  as  in  Exp.  162. 

Experiment  169  —  Fusible  Alloys 

MATERIALS.  —  Fusible  alloys,  thermometer. 

A.  Examine  specimens  of  fusible  alloys  and  state  their  charac- 
teristic properties. 

B.  Slip  a  thin  piece  of  fusible  alloy  into  a  test  tube  half  full  of  water, 
heat  the  water  gradually,  hold  a  thermometer  in  the  water,  and  note 
the  temperature  at  which  the  alloy  melts.     State  the  result. 


SODIUM  —  POTASSIUM  —  AMMONIUM 
COMPOUNDS 

Experiment  170  —  Tests  for  Sodium 

MATERIALS.  —  Sodium  compounds,  solutions  of  potassium  hydroxide 
and  tartar  emetic. 

(a)  Recall  the  flame  test,  or  apply  it  to  several  sodium  com- 
pounds, using  a  clean  test  wire  in  each  case. 

(b)  Make  a  solution  of  a  sodium  compound  slightly  alkaline 
with  potassium  hydroxide  solution,  and  add  a  little  freshly 
prepared   tartar  emetic   solution.     The  white  precipitate  is 
acid  sodium  pyroantimonate  (H2Na2Sb2O7). 

Experiment  171  —  Properties  of  Sodium  Chloride 

MATERIALS.  —  Sodium  chloride    (several  varieties)   and    the  solution 
needed  for  («). 

(a)  Examine  several  varieties  of  sodium  chloride  and  state 
the  characteristic  properties  of  each. 

(b)  Prepare  about  100  cubic  centimeters  of  a  nearly  satu- 
rated sodium  chloride  solution.     Filter,  if  it  is  not  clear,  and 
then  proceed  with  the  crystallization  as  in  Exp.  20.    Examine 
and  describe  the  best  crystals. 

(c)  Heat  a  few  crystals  of  sodium  chloride  in  a  test  tube. 
State  and  explain  the  result. 

(d)  Put  a  little  sodium  chloride  (e.g.  table  salt)  in  a  test 
tube,  and  cork  the  test  tube  tightly.    Put  some  of  the  original 
salt  in  an  open  dish.    Place  both  vessels  where  they  will  not 
be  disturbed  for  a  day  or  two,  and  then  compare  the  two  speci- 
mens.   State  and  explain  the  result. 

(e)  Apply  the  test  for  a  chloride  and  a  sulphate  to  sepa- 
rate portions  of   a  solution  of   rock   salt  and  of  table  salt. 
State  and  explain  the  results,     (i)  Test  a  specimen  of  red- 
dish rock  salt  for  iron  as  follows:    Dissolve  the  salt  in  water, 
add  a  little  dilute  hydrochloric  acid  and  boil,  cool,  and  then 


SODIUM  —  POTASSIUM  —  AMMONIUM         135 

add  ammonium  hydroxide  solution  to  alkaline  reaction;  the 
red-brown  gelatinous  precipitate  is  ferric  hydroxide.  (2)  Test  a 
specimen  of  salt  for  calcium  by  dissolving  the  solid  in  water 
and  adding  a  little  ammonium  hydroxide  and  ammonium 
oxalate  solution;  the  white  precipitate  is  calcium  oxalate. 

(3)  Test  a  specimen  of  salt  for  magnesium  as  follows:  Dissolve 
the  solid  in  water,  and  add  in  succession  ammonium  chloride 
solution,  ammonium  hydroxide,  and  disodium  phosphate  solu- 
tion;  the  white  precipitate  is  ammonium  magnesium  phos- 
phate. 

Experiment  172  —  Properties  of  Sodium  Hydroxide 
(a)  Perform,   recall,   or  repeat   (if   necessary)   experiments 
with  sodium  hydroxide  which  show  the  effect  of  (i)  exposing 
it  to  the  air,  (2)  adding  acid  to  it,  (3)  dissolving  it  in  water, 

(4)  heating  its  solution  with  aluminium  and  with  silicon. 

i  (b)  Heat  a  small  piece  of  sodium  hydroxide  upon  a  piece  of 
porcelain,  and  describe  the  result. 

(c)  Put  a  little  pulverized  sodium  hydroxide  in  a  dish  and 
let  it  stand  exposed  to  the  air  for  a  day  or  more.     Describe 
the  final  product.    Test  it  for  a  carbonate,  and  state  the  result. 

(d)  Fuse  a  small  quantity  of  sodium  hydroxide  on  a  piece 
of  porcelain,  add  a  part  of  a  match  stick  or  a  small  piece  of 
paper,  and  continue  the  fusion.    State  the  effect  on  the  wood 
and  paper. 

Experiment  173  —  Properties  of  Potassium 

MATERIALS.  —  Potassium,  litmus  paper.    . 

Precaution.      Observe   the    same    precaution   as   in   using 
sodium.     (See  Exp.  12  D.) 

(a)  Examine  a  very  small  piece  of  freshly  cut  potassium, 
and  observe  its  most  obvious  physical  properties. 

(b)  Drop  a  small  piece  of  potassium  on  the  water  in  an 
evaporating  dish.     Stand  just  near  enough  to  see  the  action. 
Describe  the  action.     Compare  it  with  the  action  of  sodium. 
Test  the  water  with  litmus  paper,  and  state  the  result.    What 
compound  of  potassium  is  in  solution? 


136  CHEMISTRY 

Experiment  174  —  Tests  for  Potassium 

MATERIALS.  —  Potassium  compounds,  sodium  cobaltinitrite  solution. 

(a)  Apply  the  flame  test  to  several  potassium  compounds, 
using  a  clean  test  wire  in  each  test.     State  the  result. 

(b)  Add  several  drops  of  sodium  cobaltinitrite  solution  to  a 
moderately  concentrated  solution  of  a  potassium  compound, 
and  shake  well.      The  yellow  precipitate  is  potassium  cobalti- 
nitrite (K3Co(NO2)6). 

Experiment  175  —  Properties  of  Ammonium  Chloride 

MATERIAL.  —  Ammonium  chloride. 

(a)  Examine  a  specimen  of  ammonium  chloride  and  state 
its  characteristic  properties. 

(b)  Add  a  few  grams  of  ammonium  chloride  to  a  test  tube 
half  full  of  water,  shake  well,  and  observe  the  result.     Does 
ammonium  chloride  dissolve  easily  in  water?    How  does  the 
dissolving  affect  the  temperature  of  the  solvent?     Save  the 
solution  for  (c). 

(c)  Add  a  small  piece  of  sodium  hydroxide  to  the  solution 
from  (b),  warm  gently,  and  very  cautiously  observe  the  odor 
of  the  gaseous  product.       What  is   the   gas?     Explain  its 
formation. 

(d)  Put  a  little  ammonium  chloride  in  a  clean,  dry  test 
tube,  heat  the  closed  end  gently,  and  observe  the  result.    What 
is  the  white  deposit?     What  general  name  is  given  to  this 
process?     To  the  product? 

SUPPLEMENTARY  EXPERIMENTS 
Experiment  176  —  Properties  of  Sodium 

MATERIALS.  —  Sodium,  litmus  paper,  tea  lead. 
Precaution.     See  Exp.  12  D. 

(a)  Examine  a  small  piece  of  sodium,  and  observe  its  most  obvious 
physical  properties,  e.g.  color,  luster,  whether  hard  or  soft. 

(b)  Perform,  recall,  or  repeat  (if  necessary)  Exp.  12  D.    (Prepara- 
tion of  Hydrogen  by  the  Interaction  of  Sodium  and  Water.) 


SODIUM  —  POTASSIUM  —  AMMONIUM         1 3  7 

(c)  Perform,  recall,  or  repeat  (if    necessary)  that  part  of    Exp. 
16  A  in  which  sodium  is  used.    (Chemical  Properties  of  Water.) 

(d)  Fill  an  evaporating  dish  nearly  full  of  water.    Put  a  piece  of 
sodium  on  a  piece  of  filter  paper  (about  the  diameter  of  the  dish), 
lay  the  paper  upon  the  water,  and  stand  back  and  observe  the  result. 
Wait  for  the  slight  explosion  that  usually  occurs  soon  after  the  action 
stops.    Describe  all  you  have  seen.    What  burned?    To  what  is  the 
vivid  color  of  the  flame  probably  due? 

Experiment  177  —  Preparation  and  Properties  of  Sodium 
Bicarbonate 

MATERIALS.  —  Ammonium  carbonate,  ammonium  hydroxide,  sodium 
chloride,  carbon  dioxide  generator. 

A.  Preparation.     Put  8  grams  of  powdered  ammonium  carbonate 
and  75  cubic  centimeters  of  ammonium  hydroxide  into  a  bottle; 
add  about  35  grams  of  fine  sodium  chloride,  cork  the  bottle,  and  shake 
the  mixture  vigorously  until  most  of  the  solid  has  dissolved.    Filter 
the  liquid,  if  it  is  not  clear,  into  a  large  test  tube.    Construct  a  carbon 
dioxide  generator  as  directed  in  Exp.  83  A.     Fill  the  generator  bottle 
half  full  of  marble,  introduce  dilute  hydrochloric  acid  as  usual,  and 
pass  carbon  dioxide  through  the  solution  from  thirty  to  forty-five 
minutes  (or  less,  if  a  precipitate  begins  to  form).    Then  remove  the 
generator,  cork  the  test  tube,  and  let  it  stand  an  hour  or  more  to 
allow  the  sodium  bicarbonate  to  settle  out  of  the  solution.    Filter, 
and  wash  quickly  with  a  very  little  cold  water.    Dry  the  precipitate 
between  filter   paper.    (Note.   If  only  a  little  of  the  precipitate  is 
formed,  use  sodium  bicarbonate  from  the  laboratory  bottle  for  B.) 

B.  Properties,     (a)  Subject  small  portions  of  the  precipitate  to 
the  flame  test  for  sodium  and  the  usual  test  for  a  carbonate.     State 
the  result. 

(b)  Put  a  little  on  moist  litmus  paper  (both  colors).    Observe  and 
explain  the  result. 

(c)  Heat  a  little  in  a  test  tube  inclined  so  that  the  open  end  is  the 
lower.    Observe  the  result.    What  is  the  visible  product?    Apply  the 
usual   test  for  carbon  dioxide  to  the  gas  in  the   test  tube;   state 
the  result.    Continue  to  heat  until  there  is  no  further  evidence  of 
change.      Determine  what  the  final  residue  is  by  applying  to  it  tests 
for  sodium,  a  bicarbonate  as  in  (a)  and  (b),  and  sodium   carbonate. 
State  the  result. 


138  CHEMISTRY 

Experiment  178  —  Testing  for  Sodium  and 
Potassium  Carbonates 

MATERIALS.  —  Washing  soda,  washing  compounds,  potash,  lye. 

Apply  the  test  for  a  carbonate,  potassium,  and  sodium  to  the 
substances  enumerated  above,  and  state  the  result  in  each  case. 

Experiment  179  —  Preparation  and  Properties  of 
Potassium  Nitrate 

MATERIALS.  —  Sodium  nitrate,  potassium  chloride,  charcoal. 

I.  Preparation.     Dissolve  about  15  grams  of  potassium  chloride 
in  about  40  cubic  centimeters  of  water,  warming  if  necessary.    Add 
about  17  grams  of  sodium  nitrate,  and  stir  well.    Boil  several  minutes, 
or  until  a  white  solid  separates.    Let  it  stand  until  the  solid  settles 
somewhat,  then  pour  the  liquid  (down  a  glass  rod  —  see  Int.  6  (i)  (a) ) 
into  an  evaporating  dish  and  let  it  cool.    Pour  off  the  liquid  from  the 
crystals.    Dissolve  the  crystals  in  a  small  volume  of  hot  water  and 
let  the  solid  crystallize  again.     Drain  off  the  water  and  dry  the 
crystals  between  filter  paper. 

II.  Properties,     (a)  Prepare  a  solution  of  the  final  crystals  and 
test  portions  for  (i)  potassium  and  (2)  a  nitrate.     State  the  result. 

0)  Test  the  solution  also  for  (i)-  sodium  and  (2)  a  chloride.  State 
the  result.  Explain  it. 

(c)  Lay  a  piece  of  charcoal  upon  a  block  of  wood  or  a  brick  and 
heat  it  by  directing  the  flame  upon  it.  Drop  potassium  nitrate 
cautiously  upon  the  hot  charcoal.  State  and  explain  the  result. 

Experiment  180  —  Properties  of  Ammonium  Compounds 

MATERIALS.  —  Ammonium  compounds. 

(a)  Recall,  perform,  or  repeat  (if  necessary)  the  experiment  show- 
ing the  effect  of  heating  ammonium  nitrate  (see  Exp.  66). 

(b)  Test  several  ammonium  salts  as  in  Exp.  175  (b)  and  (c).     State 
each  result. 

(c)  Test  baking  powder  for  ammonium  salts  (see  Exp.  113  F). 

(d)  Expose  a  piece  of  ammonium  carbonate  to  the  air.    Smell  of 
it  occasionally  and  state  the  result. 

(e)  Suggest    an    experimental   method   of   preparing   ammonium 
chloride  or  ammonium  sulphate  (see  Exp.  63).     Before  proceeding, 
submit  the  details  to  the  Teacher. 


COPPER  —  SILVER  —  GOLD 

Experiment  181  —  Properties  of  Copper 

MATERIALS.  —  Copper,  electric  bell  and  battery. 

A.  Physical,     (a)  Examine  several  forms  of  copper  —  wire, 
sheet,    filings,    borings,    etc.  —  and    state    the    characteristic 
properties. 

(b)  Hold  a  piece  of  copper  in  the  flame.    Does  it  melt  readily? 
Is  copper  a  good  conductor  of  heat?    Insert  a  piece  of  copper 
wire  in  the  circuit  with  an  electric  bell.     Is  copper  a  good 
conductor  of  electricity? 

(c)  Determine  the  specific  gravity  of  copper  (e.g.  a  compact 
roll  of  wire)  by  the  method  given  in  Exp.  88  (b).     State  the 
result. 

B.  Chemical,     (a)  Perform,  recall,  or  repeat  (if  necessary) 
the  experiments  which  show  the  effect  of  heating  copper  in 
air  (see  Exp.  4). 

(b)  Perform,  recall,  or  repeat  (if  necessary)  experiments 
which  show  the  action  of  copper  with  (i)  dilute  nitric  acid 
and  (2)  concentrated  sulphuric  acid  (see  Exps.  54,  135). 

Experiment  182  —  Tests  for  Copper 

MATERIALS.  —  Copper    wire,    copper    sulphate    solution,    ammonium 
hydroxide,  acetic  acid,  potassium  ferrocyanide  solution. 

(a)  Heat  a  copper  wire  in  the  Bunsen  flame,  and  observe 
the  color  imparted  to  the  flame.     Heat  a  minute  quantity  of 
one  or  more  copper  compounds  on  a  test  wire  in  the  flame, 
and  observe  the  color.    This  color  is  characteristic  of  copper 
and  its  compounds. 

(b)  Add  considerable  ammonium  hydroxide  to  copper  sul- 
phate solution,  shake  well,  and  observe  the  result.    The  forma- 
tion of  the  blue  solution  is  a  characteristic  and  decisive  test 
for  copper. 


140  CHEMISTRY 

(c)  Add  to  a  test  tube  one  fourth  full  of   water  an  equal 
volume  of   copper   sulphate    solution,  and    shake;    then  add 
a  few  drops  of  acetic  acid  and  of  potassium  ferrocyanide 
solution.       The    brown    precipitate    is    cupric    ferrocyanide 
(Cu2Fe(CN)6). 

(d)  Add  hydrogen  sulphide  water  to  copper  sulphate  solu- 
tion.   The  black  precipitate  is  cupric  sulphide  (CuS). 

(e)  Perform,  recall,  or  repeat  (if  necessary)  the  borax  bead 
test  for  copper. 

Experiment  183  —  Properties  of  Copper  Sulphate 

MATERIALS.  —  Copper  sulphate,  alcohol. 

(a)  Examine  a  typical  specimen  of  crystallized  copper  sul- 
phate, and  state  its  characteristic  properties. 

(b)  Prepare  anhydrous  copper  sulphate  by  heating  a  little 
of  the  pulverized  salt  in  an  evaporating  dish,     (i)  Allow  a 
little  to  remain  exposed  to  the  air  for  an  hour  or  more.    De- 
scribe and  explain  the  change  in  the  solid.     (2)   Add  the  rest 
of  the  anhydrous  copper  sulphate  to  a  test  tube  half  full  of 
alcohol,  and  shake  well.     Describe  and  explain  the  change  in 
the  solid. 

(c)  Allow  a  piece  of  red  and  of  .blue  litmus  paper  to  remain 
in  a  solution  of  copper  sulphate  for  fifteen  minutes  or  more. 
State  the  result;    explain  it  in  terms  of  the  theory  of  ioniza- 
tion.    What  term  is  applied  to  this  kind  of  a  chemical  change? 

(d)  See  Exps.  20,  22. 

Experiment  184  —  Displacement  of  Metals  —  Copper 

MATERIALS.  —  Copper  wire,  iron  nail,  zinc,  copper  sulphate  solution, 
mercuric  chloride  solution  (POISON). 

(a)  Put  a  clean  copper  wire  in  a  test  tube  half  full  of  mercuric 
chloride  solution  (POISON).     After  a  short  time  remove  the 
wire  and  wipe  it  with  a  soft  cloth  or  paper.     Observe  and 
explain  the  change. 

(b)  Put  a  clean  iron  nail  in  a  test  tube  half  full  of  copper 


COPPER  —  SILVER  —  GOLD  141 

sulphate  solution.     After  a  short  time  remove  the  nail  and 
examine  it.    What  is  the  deposit?     Explain  its  formation. 

(c)  Repeat  (b),  using  a  strip  of  zinc  instead  of  an  iron  nail. 
Observe  and  explain  the  result. 

Required  Exercise.  —  Arrange  the  metals  (used  in   this  experiment) 
in  the  order  of  their  displacing  power  with  reference  to  copper. 

Experiment  185  —  Tests  for  Silver 

MATERIALS.  —  Silver  coin,  hydrogen  sulphide,  silver  nitrate  solution. 

(a)  Recall  and  state  the  effect  of  exposing  silver  to  hydrogen 
sulphide  gas  or  to  a  sulphide  solution. 

(b)  Add  dilute  hydrochloric  acid  to  silver  nitrate  solution, 
add  considerable  ammonium  hydroxide  and  shake,  and  then 
add  dilute  nitric  acid  to  acid  reaction.    The  precipitation  of 
silver  chloride,  its  solubility  in  ammonium  hydroxide,  and  its 
reprecipitation  by  dilute  nitric  acid  constitute  the  usual  test 
for  silver. 

Experiment  186  —  Properties  of  Gold 

MATERIALS.  —  Gold,     chlorine     water,     potassium     cyanide     solution 
(POISON),  electric  bell  and  battery. 

A.  Physical,     (a)  Examine  a  specimen  of  gold  (e.g.  gold 
leaf),  and  state  its  characteristic  properties. 

(b)  Heat  a  bit  of  gold  on  charcoal  with  a  blowpipe  flame. 
Does  the  gold  melt?     Lay  a  piece  of  gold  leaf  upon  a  glass 
plate  and  touch  the  gold  with  the  two  wires  that  are  in  the 
circuit  with  an  electric  bell.    Is  gold  a  conductor  of  electricity? 

(c)  Determine  the  specific  gravity  of  a  gold  ring  by  the 
method  already  used.     State  the  result. 

B.  Chemical,     (a)  Prepare  or  obtain  about  15  cubic  centi- 
meters of  strong  chlorine  water.     Touch  a  leaf  of  gold  with 
the  moistened  end  of  a  glass  rod,  roll  the  rod  gently  over  the 
gold  to  make  some  of  the  metal  adhere,  and  lower  the  gold- 
coated  rod  carefully  into  the  chlorine  water.     Warm  gently, 
and  as  soon  as  the  gold  falls  away  from  the  rod,  remove  the 
latter  and  continue  to  warm  the  chlorine  water.     State  the 


142  CHEMISTRY 

final  result.    What  gold  compound  is  formed?    Save  the  con- 
tents of  the  test  tube  for  Exp.  187. 

(b)  Proceed  as  in  B  (a),  using  a  mixture  of  a  few  cubic  centi- 
meters of  concentrated  nitric  and  hydrochloric  acids  instead 
of  chlorine  water.    State  the  final  result.    What  gold  compound 
is  formed?     Save  the  contents  of  the  test  tube  for  Exp.  187. 

(c)  Perform  this  experiment  cautiously.     Proceed  as  in  B  (a), 
using  potassium  cyanide  solution  (POISON)  instead  of  chlorine 
water.     Heat  the  mixture  slightly.     State  the  result.     What 
gold  compound  is  formed?     Pour  the  solution  into  the  waste 
jar  in  the  hood. 

Experiment  187  —  Test  for  Gold 

MATERIALS.  —  Solutions  from  Exp.  186,  stannous  chloride  solution. 

Heat  (in  the  hood)  one  or  both  of  the  solutions  from  Exp. 
186  until  most  of  the  chlorine  has  been  driven  off,  dilute  the 
final  solution  with  water,  and  then  slowly  add  dilute  stannous 
chloride  solution.  A  precipitate  is  produced,  varying  in  color 
from  faint  purple  to  black  according  to  the  conditions.  This 
precipitate  is  finely  divided  gold;  its  formation  is  a  test  for 
gold. 

SUPPLEMENTARY  EXPERIMENTS 

Experiment  188  —  Tests  for  Copper  in  Alloys 

MATERIALS.  — •  Brass,    aluminium    bronze,    German    silver,    American 
cent,  nickel,  and  dime. 

(a)  Prepare  a  solution  of  one  of  the  alloys  enumerated  above  by 
boiling  a  small  piece  in  dilute  nitric  acid;  it  may  be  necessary  to  treat 
the  alloy  with  several  portions  of  acid  in  some  cases.    Filter  the  final 
liquid,  if  it  is  not  clear.     Apply  the  test  for  copper  to  the  clear  solu- 
tion (see  Exp.  182).    State  the  result  in  each  case. 

(b)  Proceed  as  in  (a)  with  one  or  more  alloys  obtained  from  the 
Teacher.     State  the  result. 

(c)  Proceed  as  in  (a)  with  metallic  substances  suspected  to  con- 
tain copper,  e.g.  pins  and  inexpensive  jewelry. 


COPPER  —  SILVER  —  GOLD  143 

Experiment  189  —  Preparation  and  Properties  of 

Cuprous  Oxide 

Proceed  as  in  Exp.  106.    Observe  and  state  the  properties  of  the 
precipitated  cuprous  oxide. 

,     Experiment  190  —  Deposition  of  a  Silver  Film 
Proceed  as  in  Exp.  105. 

Experiment  191  —  Displacement  of  Metals  —  Silver 
Proceed  as  in  Exp.  184,  using  silver  nitrate  solution  and  several 
metals.    State  the  result  in  each  case. 

Experiment  192  —  Tarnishing  and  Cleaning  Silver 

MATERIALS.  — •  Silver  coin,  sulphur,  rubber  band,  mustard. 

A.  Perform  one  or  more  of  the  following:   (a)  Recall  and  state  the 
effect  of  exposing  silver  to  hydrogen  sulphide  gas. 

(b)  Place  a  small  lump  of  sulphur  upon  a  clean  silver  coin,  wrap 
the  whole  tightly  in  several  pieces  of  paper,  and  let  the  package  stand 
undisturbed  for  several  days.    Examine  the  surface  of  the  coin  upon 
which  the  sulphur  was  placed.    Describe  and  explain  the  result. 

(c)  Proceed  as  in  (b),  using  a  rubber  band  instead  of  sulphur. 
Describe  and  explain  the  result. 

(d)  Cover  one  side  of  a  clean  silver  coin  with  a  paste  made  of 
mustard  and  water.    Let  the  covered  coin  stand  undisturbed  for  an 
hour  or  more.    Wash  off  the  paste,  and  examine  the  coin.    Describe 
and  explain  the  result. 

B.  Dissolve  a  little  sodium  chloride  and  sodium  bicarbonate  in 
about  75  cubic  centimeters  of  water  and  heat  to  boiling.    Put  a  small 
piece  of  aluminium  and  a  tarnished  silver  coin  into  the  solution,  taking 
care  to  have  the  metals  in  contact.    Remove  and  examine  the  coin 
after  a  few  minutes.     State  the  result. 

Experiment  193  —  Preparation  and  Properties  of  Silver 
Halides 

MATERIALS.  —  Solutions  of  silver  nitrate,  potassium  chloride,  potassium 

bromide,  potassium  iodide,  sodium  thiosulphate. 
To  separate  portions  of  silver  nitrate  solution  add  the  chloride 
bromide,  and  iodide  solution.     Observe  and  state  the  color  of  each 
precipitate.     Filter  each  separately. 


144  CHEMISTRY 

Test  precipitate  separately  by  (a)  exposing  a  little  to  the  light, 
(b)  shaking  some  with  ammonium  hydroxide,  and  (c)  shaking  some 
with  sodium  thiosulphate  solution.  State  each  result. 

Experiment  194  —  Testing  for  Copper,  Silver,  and  Gold 

(a)  Test  samples  of  inexpensive  jewelry  for  these  metals.    Cut  or 
file  the  sample  into  small  pieces  and  heat  with  dilute  nitric  acid  until 
the  solid  is  dissolved  or  there  is  no  further  evidence  of  solution. 
Filter  if  not  clear,  and  save  the  undissolved  portion,  if  any,  for  (c). 

(b)  Evaporate  the  nitrate  to  a  small  volume,  and  dilute  with 
water.     Test  this  solution  for  silver  by  adding  enough  dilute  hydro- 
chloric acid  for  complete  precipitation.    Filter,  and  test  the  nitrate 
for  copper  by  adding  ammonium  hydroxide   to   alkaline  reaction. 
State  the  result  of  each  test. 

(c)  Heat  the  undissolved  residue  from  (a)  with  aqua  regia  and  apply 
the  test  for  gold  to  the  properly  prepared  solution.     State  the  result. 


CALCIUM  —  STRONTIUM  —  BARIUM 

Experiment  195  —  Properties  of  Calcium 

MATERIALS.  —  Calcium,  electric  bell  and  battery. 

A.  Physical,     (a)  Examine  a  piece  of  clean  calcium,  and 
state  its  characteristic  properties,  e.g.  luster,  hardness. 

(b)  Insert  a  piece  of  calcium  in  the  circuit  with  an  electric 
bell.    Is  calcium  a  conductor  of  electricity? 

B.  Chemical,     (a)  Let   a   piece   of   clean   calcium   remain 
exposed  to  the  air  for  several  days.    Describe  the  final  result. 

(b)  Heat  a  test  tube  half  full  of  water,  nearly  to  boiling,  and 
drop  in  several  small  pieces  of  calcium.    Observe  and  describe 
the  action.    Test  the  gas  with  a  blazing  joss  stick.    What  is 
the  gas?     Describe  the  contents  of  the  test  tube.    What  is  the 
suspended  solid?     Write  the  equation  for  the  interaction  of 
calcium  and  water. 

(c)  Heat  a  small  piece  of  calcium  several  minutes  on  char- 
coal in  the  oxidizing  flame  of  a  blowpipe.     State  the  result. 
What  is  the  product? 

(d)  Drop  a  small  piece  of  calcium  into  a  test  tube  one 
fourth  full  of  dilute  hydrochloric  acid,  and  warm  gently  if 
the  action  is  not  marked.    State  the  result.     If  a  gas  is  liber- 
ated, test  it  with  a  lighted  joss  stick;  state  the  result.      Write 
the  equation  for  the  interaction  of  calcium  and  hydrochloric 
acid. 

(e)  Proceed  as  in  (d),  using  dilute  nitric  acid  instead  of 
hydrochloric  acid. 

(/)  Proceed  as  in   (d),  using  dilute  sulphuric  acid  instead 
of  hydrochloric  acid. 


146  CHEMISTRY 

Experiment  196  —  Tests  for  Calcium 

MATERIALS.  —  Calcium  compounds,  ammonium  oxalate  and  ammonium 
carbonate  solutions. 

(a)  Subject  several  calcium  compounds  to  the  flame  test, 
using  a  clean  test  wire  in  each  case.     What  color  is  imparted 
to  the  flame? 

(b)  Add  an  excess  of  ammonium  oxalate  solution  to  calcium 
chloride  solution,  and  state  the  result.     The  precipitate  is 
calcium  oxalate.    Divide  into  two  parts.    To  (i)  add  an  excess 
of  dilute  hydrochloric  acid,  warm  gently,  and  state  the  final 
result.    To  (2)  add  considerable  acetic  acid  and  warm  gently; 
observe  and  state  the  final  result. 

(c)  Add  an  excess  of  ammonium  carbonate  solution  to  cal- 
cium chloride  solution,  and  state  the  result.    The  precipitate  is 
calcium  carbonate.     Divide  it  into  two  parts,  and  treat  with 
the  acids  as  in  (b).    State  the  results  and  compare  with  the 
results  obtained  in  (b). 

(d)  Suggest  a  test  for  calcium  in  calcium  carbonate  and 
calcium  sulphate. 

Experiment  197  —  Testing  for  Calcium 

MATERIALS.  —  Mortar,    plaster,    bone    ash,    plaster    of    Paris,    tooth 
powder,  whiting,  cement,  bleaching  powder. 

(a)  Prepare  a  solution  of  the  substances  enumerated  above 
by  boiling  a  little  of  each  with  dilute  hydrochloric  acid  (or 
dilute  nitric  acid)  and  filtering.    Test  the  filtrate  for  calcium. 
State  the  result  in  each  case. 

(b)  Obtain  "unknowns"  from  the  Teacher  and  test  them 
for  calcium. 

Experiment  198" —  Plaster  of  Paris 

MATERIALS.  —  Plaster  of  Paris,  block  of  wood,  coin,  vaseline. 

Mix  a  little  plaster  of  Paris  with  enough  water  on  a  block 
of  wood  to  form  a  thick  paste.  Rub  a  very  little  vaseline  upon 
one  side  of  a  coin,  and  press  the  coin,  coated  side  down,  into  the 
paste.  Let  it  stand  undisturbed  for  fifteen  or  more  minutes. 


CALCIUM  —  STRONTIUM  —  BARIUM  147 

Then  remove  the  coin  carefully,  and  examine  and  describe 
the  effect  upon  the  hardened  plaster. 

Experiment  199  —  Calcium  Compounds  and 
Hardness  of  Water    ' 

(a)  Proceed  as  in   Exp.    116   (d),  using  only   the   calcium 
compounds. 

(b)  Prepare  some  permanently  hard  water  and  devise  an 
experiment  to  soften  it.     Submit  the  details  to  the  Teacher 
before  proceeding. 

Experiment  200  —  Tests  for  Strontium 

MATERIALS.  —  Strontium    compounds,    test    wire,    calcium    sulphate 

solution. 

(a)  Apply  the  flame  test  to  strontium  nitrate  and  other 
available  strontium  compounds,   using  a  clean  test  wire  in 
each  case.     What  color  is  imparted  to  the  flame?     Compare 
this  color  with  that  produced  by  calcium  compounds. 

(b)  To  the  solution  of  a  strontium  compound  add  calcium 
sulphate  solution.     The  precipitate  is  strontium  sulphate. 

Experiment  201  —  Tests  for  Barium 

MATERIALS.  —  Barium   compounds,   test  wire,   potassium  dichromate 

solution. 

(a)  Apply  the  flame  test  to  barium  nitrate  and  other  avail- 
able barium  compounds,  using  a  clean  test  wire  in  each  case. 
What  color  is  imparted  to  the  flame?     Compare  this  color 
with  that  produced  by  calcium  and  by  strontium  compounds. 

(b)  Add  dilute  sulphuric  acid  to  barium  chloride  solution 
(or  the  solution  of   any  barium  compound).     The  precipitate 
is  barium  sulphate.    Describe  it.    Test  its  solubility  by  heat- 
ing a  little  of  the  precipitate  in  (i)  concentrated  hydrochloric 
acid,  (2)  concentrated  nitric  acid,  (3)  concentrated  sulphuric 
acid;   perform  the  experiment  in  the  hood  and  heat  the  acids 
cautiously,  especially  the  sulphuric  acid.      State  the  results. 

(c)  Add  potassium  dichromate  solution  to  barium  nitrate 


148  CHEMISTRY 

solution.  The  precipitate  is  barium  chromate.  Describe  it. 
Test  its  solubility  by  heating  some  of  the  precipitate  in  (i) 
acetic  acid  and  (2)  concentrated  hydrochloric  acid.  State 
the  results. 

SUPPLEMENTARY  EXPERIMENTS 

Experiment  202  —  Calcium  Carbonate  and  Acid 
Calcium  Carbonate 

Perform,  recall,  or  repeat  (if  necessary)  the  experiment  in  which 
gentle  heat  was  applied  to  the  product  of  the  interaction  of  an  excess 
of  carbon  dioxide  and  calcium  hydroxide.  (See  Exp.  85.)  Express 
the  essential  chemical  changes  by  reactions. 

Experiment  203  —  Preparation  and  Properties  of  Calcium 
Oxide  and  Calcium  Hydroxide 

MATERIALS.  —  Calcium  carbonate,  calcium  oxide. 

A.  Preparation,     (a)  Wind  a  test  wire  around  a  small  lump  of  cal- 
cium carbonate,  and  heat  the  solid  for  several  minutes  in  the  hottest 
part  of  the  Bunsen  flame;  or  heat  the  calcium  carbonate  on  charcoal 
with  the  oxidizing  flame  of  the  blowpipe.     Then  let   the  residue 
cool  somewhat,  put  it  in  an  evaporating  dish,  and  add  a  little  water. 
Observe  the  result.    Test  the  liquid  with  red  litmus  paper;    test  it 
also  for  calcium.     State  the  results.     What  calcium  compound  was 
formed  by  heating  the  calcium  carbonate?    By  treating  the  product 
of  the  heating  with  water? 

(b)  Prepare  a  small  quantity  of  solid  calcium  hydroxide  by  adding 
a  little  water  to  a  lump  of  lime,  and  save  it  for  C. 

B.  Properties  of  Calcium  Oxide,     (a)  Examine  a  lump  of  calcium 
oxide  and  state  its  characteristic  properties. 

(b)  Put  a  lump  of  calcium  oxide  on  a  glass  plate  or  block  of  wood 
and  let  it  remain  exposed  to  the  air  for  a  few  days.    Examine  it  at 
intervals  and  describe  it.    Describe  the  final  product.    What  is  it? 

(c)  Recall,  perform,  or  repeat  (if  necessary)  the  experiment  that 
shows  the  effect  of  mixing  calcium  oxide  and  water.     Express  the 
chemical  reaction  by  an  equation. 

C.  Properties  of  Calcium  Hydroxide,     (a)  Examine  the  calcium 
hydroxide  prepared  in  A  (ft)  and  state  its  characteristic  properties. 


CALCIUM  —  STRONTIUM  —  BARIUM  149 

(b)  Add  a  little  calcium  hydroxide  to  a  test  tube  half  full  of  water 
and  shake  vigorously.    Let  the  suspended  solid  settle  somewhat,  and 
filter.    Pour  half  of  the  filtrate  into  an  evaporating  dish  and  evaporate 
it  to  dryness.     (Meanwhile  (c)  may  be  performed.)     Compare  the 
amount  of  residue  in  the  dish  with  the  amount  originally  shaken  with 
water.    Draw  a  conclusion  regarding  the  solubility  of  calcium  hydrox- 
ide in  water. 

(c)  Taste  of  the  solution  saved  from  (b),  and  describe  the  taste. 
Determine  the  reaction  toward  litmus;  is  the  solution  acid,  alkaline, 
or  neutral?    Heat  the  solution  slowly  to  boiling,  and  describe  the 
result.     What  is  the  effect  of  increased  heat  on  the  solubility  of  cal- 
cium hydroxide  in  water? 

(d)  State  the  result  of  (i)  exposing  calcium  hydroxide  solution  to 
the  air  and  (2)  exhaling  the  breath  through  calcium  hydroxide  solu- 
tion.    Express  each  reaction  by  an  equation. 

Experiment  204  —  Preparation  of  Red  Fire  and  Green  Fire 

MATERIALS.  —  Strontium  nitrate,  powdered  potassium  chlorate,  pow- 
dered shellac,  iron  pan  or  brick,  barium  nitrate. 

A.  Mix  carefully  small  and  equal  (in  bulk)  quantities  of  the  three 
substances  on  a  sheet  of  paper.    Place  the  mixture  in  an  iron  pan  or 
on  a  brick  in  the  hood,  and  light  it  with  a  Bunsen  burner.    Describe 
the  result. 

B.  Proceed  as  in  A,  using  barium  nitrate  instead  of  strontium 
nitrate. 


ALUMINIUM 

Experiment  205  —  Properties  of  Aluminium 

A.  Physical.     Proceed  as  in  Exp.  181  A  (a),  (b),  (c),  using 
aluminium  instead  of  copper. 

B.  Chemical,     (a)  Warm  a  piece  of  aluminium  with  con- 
centrated hydrochloric  acid.     Test  the  escaping  gas  with  a 
blazing  joss  stick;    what  is  the  gas?     What  compound  of 
aluminium  is  formed? 

(b)  Boil  a  piece  of  aluminium  with  concentrated  sodium 
hydroxide  solution.  Test  as  in  B  (a).  What  is  the  gas?  What 
compound  of  aluminium  is  formed? 

Experiment  206  —  Preparation  and  Properties  of 
Aluminium  Hydroxide 

MATERIALS.  —  Solutions  of    aluminium   sulphate,   sodium    hydroxide, 
potassium  hydroxide,  ammonium  sulphide,  and  sodium  carbonate. 

A.  Preparation,     (a)  Add  ammonium  hydroxide  to  a  solu- 
tion of  aluminium  sulphate,  and  shake  well.    The  precipitate 
is  aluminium  hydroxide;  save  it  for  further  use  in  this  experi- 
ment. 

(b)  Proceed  as  in  (a),  using  aluminium  sulphate  solution 
and  a  very  little  sodium  hydroxide  solution.  Compare  with 
the  result  in  (a).  Save  this  precipitate.  Predict  the  result 
of  using  potassium  hydroxide  (instead  of  sodium  hydroxide). 
Verify  the  prediction. 

B.  Properties,     (a)  Examine   the   precipitate   from   A    (a) 
and  note  its  properties,  e.g.  color,  texture,  etc.     Remove  a 
little  and  rub  it  between  the  fingers;   describe  the  result. 

(b)  To  the  precipitate  from  A  (b)  add  sodium  hydroxide 
slowly  and  shake  constantly  until  a  conspicuous  change  occurs. 
State  the  result.  What  compound  of  aluminium  is  formed? 


ALUMINIUM  151 

(c)  To  a  portion  of  the  precipitate  from  A  (a)  add  consider- 
able ammonium  hydroxide,  and  shake  well.     Compare  with 
the  result  in  B  (b). 

(d)  Add  dilute  hydrochloric  acid  to  a  portion  of  the  precipi- 
tate from  A  (a),  and  shake  well.     State  the  result.     Proceed 
similarly  with  other  acids,  e.g.  sulphuric  and  acetic.     State 
the  results. 

Experiment  207  —  Clarification  of  Water  by 
Aluminium  Hydroxide 

MATERIALS.  —  Clay,  aluminium  sulphate  solution. 

Shake  a  little  fine  clay  (or  clay  soil)  in  about  50  cubic  centi- 
meters of  water.  Divide  the  turbid  water  into  two  equal 
portions.  Save  one  for  comparison.  To  the  other  add  a  little 
aluminium  sulphate  solution  and  ammonium  hydroxide,  and 
mix  well.  Compare  the  two  portions  in  a  few  minutes.  State 
the  final  result.  Explain  it. 

Experiment  208  —  Thermit 

MATERIALS.  —  Thermit,  sand  (or  clay)  crucible  about  10  centimeters 
(4  inches)  deep,  brick,  sand,  iron  pan,  granulated  aluminium, 
barium  peroxide,  magnesium  ribbon. 

Fill  an  iron  pan  with  sand  and  stand  it  on  a  brick.  Put  about 
30  grams  of  thermit  in  the  crucible,  and  bury  the  crucible 
about  halfway  in  the  sand;  as  a  precaution,  a  box  of  sand 
should  be  near  by  to  throw  upon  the  molten  iron  in  case  the 
crucible  should  break.  Prepare  a  fuse  mixture  by  mixing  thor- 
oughly about  10  grams  of  barium  peroxide  and  i  gram  of 
granulated  aluminium.  Make  a  hole  in  the  top  of  the  thermit 
and  pour  in  the  fuse  mixture;  insert  a  piece  of  magnesium 
ribbon  into  the  heap  of  fuse  mixture.  Light  the  magnesium 
with  the  Bunsen  flame,  and  stand  away.  The  reaction  is 
vigorous.  Describe  it.  When  the  crucible  is  cool,  break  it 
open  and  examine  the  contents.  Describe  the  two  parts. 
What  is  the  name  of  each? 


152  CHEMISTRY 

Experiment  209  —  Tests  for  Aluminium 

MATERIALS.  —  Aluminium  sulphate,  cobaltous  nitrate  solution,  blow- 
pipe,   charcoal. 

(a)  Proceed  as  in  Exp.  206  A  (a). 

(b)  To  a  portion  of  the  solution  of  the  aluminium  compound 
add  a  little  sodium  hydroxide  solution  and  then  an  excess. 
To  another  portion  add  an  excess  of  ammonium  hydroxide. 
The   precipitation   and   properties   of   aluminium   hydroxide 
serve  as  the  test. 

(c)  Heat  a  little  aluminium  sulphate  (or  any  other  alumin- 
ium compound)   on  charcoal  in  the  blowpipe  flame.     Cool, 
and  moisten  with  a  drop  or  two  of  cobaltous  nitrate  solution. 
Heat  again,  and  observe  the  color  of  the  residue. 

SUPPLEMENTARY  EXPERIMENTS 

Experiment  210  —  Aluminium  Salts  as  Mordants 

MATERIALS.  —  Solutions  of  alum  and  cochineal. 

Add  a  little  alum  solution  to  a  dilute  solution  of  cochineal,  then  add 
ammonium  hydroxide  and  shake  well.  Filter,  and  compare  the  colors 
of  the  nitrate  and  precipitate. 

Experiment  211  —  Preparation  and  Properties  of 
Potassium  Alum 

MATERIALS.  —  Aluminium  sulphate,  potassium  sulphate,  evaporating 

dish. 

A.  Preparation.    Add  8  grams  of  aluminium  sulphate  and  4  grams 
of  potassium  sulphate  to  about  40  cubic  centimeters  of  water  and  heat 
the  mixture  in  a  porcelain  dish*  until  the  salt  is  dissolved.     Set  it  aside 
to  crystallize;  well-formed  crystals  may  be  obtained  upon  a  thread 
suspended  in  the  solution.    (Meanwhile  proceed  with  B  (a)) .    Crystals 
of  potassium  alum  will  be  deposited.    Remove  the  best  ones;  dry  and 
examine.    Describe  them,  giving  color,  luster,  size,  and  crystal  form. 

B.  Properties,     (a)  Test  a  solution  of  alum  (from  the  laboratory 
bottle)  for  aluminium  ions  and  sulphate  ions,  and  state  the  result. 
Cautiously  taste  the  solution,  and  describe  the  result.     Test  solid 
alum  for  water  of  crystallization,  and  state  the  result. 


ALUMINIUM  153 

(6)  Select  several  good  crystals  from  those  prepared  in  A  and  ex- 
amine them  carefully.  Describe  them.  Test  as  in  B  (a),  and 
state  the  results.  Allow  some  crystals  to  remain  exposed  to  the  air 
for  several  hours.  Compare  finally  with  the  original  crystals. 
Explain  the  difference. 

Experiment  212  —  Displacement  of  Metals  by  Aluminium 
Devise  experiments  similar  to  Exp.  184  to  illustrate  the  displace- 
ment of  metals  by  aluminium. 

Experiment  213  —  Equivalent  of  Aluminium 

Proceed  as  in  Exp.  64. 

Experiment  214  —  Hydrolysis  of  Aluminium  Salts 

(a)  Prepare  a  solution  of  aluminium  sulphate  and  test  it  with  lit- 
mus paper.     State  and  explain  the  result. 

(b)  Proceed  as  in  (a)  with  an  alum  solution.     State  and  explain 
the  result. 

Experiment  215  —  Alum  Baking  Powder 
Proceed  as  in  Exp.  113  G. 


IRON 

Experiment  216  —  Properties  of  Iron 

MATERIALS.  —  Cast  and  wrought  iron,  steel,  magnet,  iron  thread,  iron 

powder. 

A.  Physical,     (a)  Examine  typical  specimens  of  cast  iron, 
wrought  iron,  and  steel,  and  state  their  characteristic  physical 
properties. 

(b)  Determine  the  heat  and  the  electrical  conductivity  of 
iron  wire  by  proceeding  as  in  Exp.  181   (b).     Compare  the 
result  with  that  of  similar  experiments. 

(c)  Determine  the  specific  gravity  of  iron  by  the  method 
given  in  Exp.  88  (b).    Compare  as  in  (b). 

(d)  Try  the  action  of  a  magnet  on  each  kind  of  iron.    State 
the  result. 

(e)  Drop  a  pinch  of  iron  powder  into  a  Bunsen  flame.    Hold 
a   piece  of   iron   thread  in  the  flame.     Describe  the  results, 
and  draw  a  conclusion. 

B.  Chemical,     (a)  Perform,  recall,  or  repeat  (if  necessary) 
experiments  which  show  the  effect  of  heating  iron  in  oxygen, 
chlorine,   nitrogen,   nitrous  oxide,   and  sulphur.     State  each 
result. 

(b)  As  in  (a),  experiments  showing  the  action  of  acids  with 
iron.     State  the  results. 

(c)  As  in  (a)  experiments  illustrating  the  displacement  of 
metals  by  iron.    If  necessary,  try  additional  experiments  with 
iron  and  solutions  of  metals.     State  the  results.     (Compare 
Exps.  184,  191,  212.) 

Experiment  217  —  Properties  of  Ferrous  Compounds 

MATERIALS.  —  Iron  powder  (or  filings),  hydrochloric  acid,  solutions  of 
sodium  hydroxide  and  potassium  ferricyanide. 

Put  a  few  grams  of  iron  powder  in  a  test  tube,  add  about 
10  cubic  centimeters  of  dilute  hydrochloric  acid,  and  warm 


IRON  155 

gently;  ferrous  chloride  is  formed  (in  solution).  Proceed  as 
follows:  (i)  Pour  a  little  into  a  test  tube  one  third  full  of 
sodium  hydroxide  solution.  The  precipitate  is  ferrous  hydrox- 
ide. Describe  it.  Watch  and  describe  the  changes  in  color. 
To  what  are  the  changes  due?  (2)  Add  a  second  portion  to 
potassium  ferricyanide  solution.  The  precipitate  is  ferrous 
ferricyanide.  Describe  it. 

Experiment  218  —  Properties  of  Ferric  Compounds 

MATERIALS.  —  Solutions  of  ferric  chloride,  sodium  hydroxide,  potassium 
sulphocyanate,  and  potassium  ferrocyanide. 

To  a  little  ferric  chloride  solution  add  (i)  sodium  hydroxide 
solution.  The  precipitate  is  ferric  hydroxide.  Describe  it. 
Add  to  ferric  chloride  solution  (2)  a  little  potassium  sulpho- 
cyanate solution.  The  rich  wine-red  color  is  caused  by  soluble 
ferric  sulphocyanate.  This  test  readily  distinguishes  'ferric 
from  ferrous  compounds.  Add  as  above  (3)  a  little  potassium 
ferrocyanide  solution.  The  precipitate  is  ferric  ferrocyanide. 
Describe  it. 

Tabulate  the  results  of  Exps.  217  and  218,  showing  the  be- 
havior of  ferrous  and  ferric  compounds  under  the  same 
conditions. 

Experiment  219  —  Interrelation  of  Iron  Compounds 

MATERIALS.  —  Ferric  chloride  solution,  zinc,  ferrous  sulphate,  potas- 
sium chlorate. 

A.  Put  a  piece  of  zinc  in  ferric  chloride  solution  made  slightly 
acid  by  hydrochloric  acid.     After  the  operation  has  proceeded 
for  about  fifteen  minutes,  test  a  portion  of  the  liquid  for  a 
ferrous  and  a  ferric  compound  by  Exps.  217  and  218.     State 
and  explain  the  result. 

B.  (a)  To  a  solution  prepared  from  fresh  or  freshly  washed 
ferrous  sulphate  add  a  little  hydrochloric  acid,  warm  gently, 
and  then  add  a  few  crystals  of  potassium  chlorate.    After  heat- 
ing a  short  time,  test  portions  of  the  solution  'for  a  ferric  and 
a  ferrous  compound  (as  in  A).     State  and  explain  the  result. 


156  CHEMISTRY 

(b)  Add  10  cubic  centimeters  of  concentrated  nitric  acid, 
drop  by  drop,  to  a  hot  solution  of  ferrous  sulphate  to  which  a 
little  sulphuric  acid  has  been  added,  and  boil.  Test  and  explain 
as  in  (a). 

SUPPLEMENTARY  EXPERIMENTS 

Experiment  220  —  Testing  Substances  for  Iron 

MATERIALS.  —  Clay,  brick,  flower  pot,  bauxite,  rusty  rock,  sheet  tin, 
iron  rust,  bluing,  solutions  of  potassium  ferricyanide  and  potassium 
sulphocyanate. 

(a)  Prepare  a  solution  of  each  of  the  substances  to  be  tested  by 
boiling  a  little  of  the  powdered  material  with  concentrated  hydro- 
chloric acid.    Test  the  clear  solution  for  iron,  both  ferric  and  ferrous, 
and  state  the  result  in  each  case. 

(b)  Obtain  other  substances  (from  the  Teacher)  and  test  them  for 
iron  as  in  (a). 

(c)  Add  considerable  sodium  hydroxide  solution  to  a  dilute  solu- 
tion  of  bluing.     Describe  and  explain   the   result.    (SUGGESTION. 

-See  Exp.  218  (i).) 

Experiment  221  —  Blue  Print  Paper 

MATERIALS.  —  Ferric  ammonium  citrate  solution  (2.2  gm.  to  10  cc.  of 
water),  potassium  ferricyanide  solution  (same  concentration), 
cotton  or  sponge,  smooth,  hard  paper. 

Prepare,  or  obtain  from  the  Teacher,  the  two  solutions.  Mix 
equal  volumes  of  each.  Dip  a  piece  of  absorbent  cotton  or  a  sponge 
into  the  mixture,  squeeze  out  the  excess,  and  apply  a  thin  coat  to  one 
side  of  a  piece  of  paper.  Let  the  paper  dry  in  a  dark  place. 

Lay  a  key  or  a  coin  on  the  coated  paper  and  expose  to  the  sunlight 
for  a  few  minutes.  Wash  the  paper  thoroughly  with  water.  Com- 
pare it  with  (i)  the  original  uncoated  paper,  (2)  the  coated  unexposed 
paper,  and  (3)  a  piece  of  exposed  and  washed  paper.  State  the 
result. 


IRON  157 

Nickel  and  Cobalt 

Experiment  222  —  Test  for  Nickel 

To  a  solution  of  nickel  chloride  add  sodium  hydroxide  to  alkaline 
reaction.  The  precipitate  is  nickelous  hydroxide.  Describe  it. 

Experiment  223  —  Test  for  Cobalt 

To  a  solution  of  cobaltous  nitrate  add  acetic  acid  and  considerable 
solid  potassium  nitrite,  warm  and  shake  well.  A  yellow  precipitate 
of  potassium  cobaltinitrite  (K3Co(NO2)6)  is  formed. 


MAGNESIUM  —  ZINC  —  MERCURY  —  CADMIUM 

Experiment  224  —  Tests  for  Magnesium 

MATERIALS.  —  Solutions   of    magnesium    sulphate    (or    chloride),    am- 
monium chloride,  ammonium  hydroxide,  disodium  phosphate. 

(a)  To  a  solution  of  magnesium  sulphate  (or  chloride)  add 
successively    solutions  of    ammonium    chloride,    ammonium 
hydroxide,  and  disodium  phosphate.     A  precipitate  of  ammo- 
nium  magnesium  phosphate  (NHUMgPOO   is  formed.     De- 
scribe it. 

(b)  Put  a  little  powdered  magnesium  oxide  in  a  cavity  at 
the  end  of  a  piece  of  charcoal,  moisten  with  water,  and  heat 
intensely  in  a  blowpipe  flame.       Cool,  and  moisten  with  a 
drop  of  cobaltous  nitrate  solution.     Heat  again,  and  when 
cool  observe  the  color.    If  the  experiment  has  been  conducted 
properly,  a  pink  or  pale  flesh  colored  residue  coats  the  charcoal. 

Experiment  225  —  Tests  for  Zinc 

MATERIALS.  —  Zinc,  zinc  sulphate  and  cobalt  nitrate  solutions,  zinc 
oxide,  blowpipe,  .charcoal. 

(a)  Apply  tests  for  metallic  zinc  (see  Exp.  181,  A). 

(b)  Add  a  very  little  sodium  hydroxide  solution  to  zinc 
sulphate  solution,  and  shake  well.     Describe  the  precipitate. 
What  is  it?    Divide  the  mixture  into  three  parts.     To  one  add 
considerable  sodium  hydroxide,  to   another  add   ammonium 
hydroxide,   and  to  the  third  add  dilute  hydrochloric  acid. 
Shake  each  well,  and  observe  the  result.      What  compound  of 
zinc  is  formed  in  each  case? 

(c)  Fill  a  small  cavity  at  one  end  of  a  piece  of  charcoal 
with  zinc  oxide   (or   any  other  zinc   compound),   and  heat 
intensely  in  the  blowpipe  flame.      Moisten  with  a  drop  of 
cobaltous  nitrate  solution,  then  heat  again.    Cool  and  examine. 
State  the  color  of  the  incrustation  on  the  charcoal  in  or  near 
the  cavity.     (Compare  Exp.  209  (c).) 


MAGNESIUM — ZINC — MERCURY — CADMIUM  1 59 

Experiment  226  —  Properties  of  Mercurous  and  Mercuric 
Compounds  and  Tests  for  Combined  Mercury 

MATERIALS.  —  Solutions  of  mercurous  nitrate,  mercuric  chloride,  and 
potassium  iodide. 

A.  Mercurous.     (a)  Add  a  few  drops  of  hydrochloric  acid 
to  a  little  mercurous  nitrate  solution.    The  precipitate  is  mer- 
curous chloride.     Describe  it.     Note  its  insolubility  in  water 
and  in  dilute  hydrochloric  acid.    Add  an  excess  of  ammonium 
hydroxide.     The  black  precipitate  is  a  test  for  mercury  in 
mercurous  compounds. 

(b)  Add  potassium  iodide  solution  to  mercurous  nitrate 
solution.  Describe  the  result.  What  compound  of  mercury 
is  formed? 

B.  Mercuric,     (a)  Add  a  few  drops  of  hydrochloric  acid  to 
a  little  mercuric  nitrate  solution.     Compare  with  the  result  in 
A  (a).    Add  a  few  drops  of  ammonium  hydroxide,  or  enough 
to  produce  a  decided  change.    Compare  with  A  (a).    The  pre- 
cipitate is  mercuric  ammonium  chloride. 

(b)  Proceed  as  in  A  (b),  using  mercuric  chloride  solution. 

C.  See   experiments   showing     displacement   of   metals   in 
which  mercury  and  its  compounds  are  involved. 

SUPPLEMENTARY  EXPERIMENTS 

Experiment  227  —  Properties  of  Magnesium  and  Zinc 

MATERIALS.  —  Magnesium,  zinc. 

A.  Physical.     Proceed  as  in  Exp.  181  A,  using  magnesium  and 
zinc  instead  of  copper. 

B.  Chemical,     (a)  Perform,  recall,  or  repeat  (if  necessary)  experi- 
ments showing  the  results  of  heating  magnesium  and  zinc  (i)  in  a 
limited  supply  of  air  and  (2)  in  an  abundance  of  air;  and  also  treat- 
ing magnesium  and  zinc  with  acids.     State  the  results. 

(b)  Proceed  as  in  Exp.  205  B,  using  zinc  instead  of  aluminium. 
What  zinc  compound  is  formed? 

Experiment  228  —  Equivalent  of  Zinc  and  Magnesium 
Proceed  as  in  Exps.  61  and  62. 


160  CHEMISTRY 

Experiment  229  —  Physical  Properties  of  Mercury 

(a)  Pour  a  drop  or  two  of   mercury  into  an  evaporating  dish. 
Examine  the  mercury,  and  state  its  characteristic  properties.    Agi- 
tate the  dish,  and  describe  the  result. 

(b)  Stand  a  funnel  in  a  test  tube  and  carefully  pour  the  mercury 
from  the  dish  into  the  test  tube.    Remove  the  funnel.    Heat  the  bot- 
tom of  the  test  tube  gently  and  observe  the  result,  especially  the 
deposit,  if  any,  upon  the  upper  part  of  the  tube.     Scrape  a  little  of 
it  out  of  the  tube  with  a  glass  rod.    What  is  the  deposit?    What  prop- 
erty of  mercury  is  shown  by  this  experiment? 

(c)  Suggest   a   method   of   determining    the   specific   gravity   of 
mercury.    If  approved  by  the  Teacher,  try  it. 

Experiment  230  —  Displacement  of  Metals  —  Magnesium, 

Zinc,  and  Mercury 

Proceed  as  in  former  experiments  (e.g.  Exp.  184),  using  these  metals 
and  solutions  of  several  metallic  compounds.     Tabulate  the  results. 

Experiment  231  —  Test  for  Cadmium 

MATERIALS.  —  Solutions  of  cadmium  chloride  and  hydrogen  sulphide. 

Add  hydrogen  sulphide  water  to  a  test  tube  half  full  of  cadmium 

chloride  solution.    The  precipitate  is  cadmium  sulphide.    Describe  it. 


TIN   AND    LEAD 

Experiment  232  —  Test  for  Tin 

MATERIALS.  —  Solutions  of  stannous  chloride  and  mercuric  chloride 

(POISON). 

Add  a  few  drops  of  stannous  chloride  solution  to  mercuric 
chloride  solution.  The  white  precipitate  is  mercurous  chloride. 
Add  considerable  stannous  chloride  and  warm  gently.  The 
gray-black  precipitate  is  finely  divided  mercury.  Explain  the 
chemical  change  and  express  it  by  equations. 

,:• 

Experiment  233  —  Tests  for  Lead 

MATERIALS.  —  Lead    nitrate    and    potassium    dichromate    solutions, 
sulphuric  acid,  hydrochloric  acid. 

(a)  Recall   the  result  of  reducing  lead  oxide  in  the  blow- 
pipe flame. 

(b)  Recall  the  action  of  hydrogen  sulphide  with  the  solution 
of  a  lead  compound. 

(c)  Add  dilute  hydrochloric  acid  to  a  little  lead  nitrate  solu- 
tion until  precipitation  ceases.     The  insoluble  precipitate  is 
lead  chloride.    Boil  some  of  the  precipitate  with  considerable 
water.     Describe  the  action. 

(d)  Add  dilute  sulphuric  acid  to  a  little  lead  nitrate  solution 
until  precipitation  ceases.     The  precipitate  is  lead  sulphate. 
Observe  its  properties.     Is  it  soluble  in  hot  water?     Try  it. 

(e)  Repeat   (d),   using  potassium  dichromate  solution  in- 
stead of  sulphuric  acid.     The  precipitate  is  lead  chromate. 
Describe  it,  especially  the  color. 

SUPPLEMENTARY  EXPERIMENTS 

Experiment  234  —  Physical  Properties  of  Tin  and  Lead 

MATERIALS.  —  Tin,  lead. 

(a)  Proceed  with  tin  and  lead  as  in  former  experiments  (e.g. 
Exp.  181). 


1 62  CHEMISTRY 

(b)  Bend  a  piece  of  tin  and  note  the  crackling  noise. 

(c)  Rub  a  piece  of  lead  upon  a  hard  surface,  e.g.  the  (outside) 
bottom  of  a  mortar.    State  the  result.    Rub  lead  with  the  fingers, 
and  compare  with  the  preceding  result. 

Experiment  235  —  Displacement  of  Metals  —  Tin  and  Lead 

Proceed  as  in  former  experiments  (e.g.  Exp.  184),  using  tin  and 
lead  and  solutions  of  metals.  Tabulate  the  results  and  compare  with 
similar  experiments. 

Experiment  236  —  Testing  for  Lead  and  Tin 

MATERIALS.  —  As  below. 

A.  Lead,     (a)  Warm  thin  shavings  of  solder  with  dilute  nitric 
acid,  filter  and  test  the  filtrate  for  lead. 

(b)  Proceed  as  in  (a)  using  one  or  more  of  the  following:  Tea  lead, 
type  metal,  plumbago,  shot,  bullets,  metallic  cap  of  a  bottle,  "lead" 
of  a  lead  pencil,  and  "unknowns"  obtained  from  the  Teacher. 

(c)  Apply  the  (reduction)  blowpipe  test  for  lead  to  white  lead, 
red  lead,  litharge,  dry  paint,  chrome  yellow,  and  "unknowns." 

B.  Tin.     (a)  Boil  a  piece  of  tin  foil  with  concentrated  hydro- 
chloric acid,  filter,  and  test  the  filtrate  for  tin. 

(b)  Proceed  as  in  B  (a)  with  metals  known  or  suspected  to  contain 
tin. 

Experiment  237  —  Analysis  of  Solder 
MATERIALS.  —  Solder,  ammonium  polysulphide  solution. 

Dissolve  a  gram  of  solder  filings  in  as  small  a  quantity  of  hot 
aqua  regia  as  possible,  evaporate  nearly  to  dryness  (in  the  hood), 
dissolve  the  residue  in  water,  add  10  to  15  cubic  centimeters  of 
dilute  hydrochloric  acid,  and  precipitate  the  metals  as  sulphides  by 
bubbling  hydrogen  sulphide  gas  through  the  solution  for  15  or  20 
minutes.  Filter,  wash  with  hot  water,  pierce  a  hole  in  the  filter 
paper,  and  wash  the  precipitate  into  a  test  tube  with  yellow  am- 
monium sulphide.  Add  more  ammonium  sulphide,  and  shake. 
Filter. 

The  filtrate  contains  the  tin  as  ammonium  sulphostannate;  add 
to  it  dilute  hydrochloric  acid,  and  yellow  stannic  sulphide  is 
precipitated. 

The  precipitate  is  lead  sulphide.  Dissolve  it  in  hot  dilute  nitric 
acid,  filter,  and  test  the  filtrate  for  lead. 


TIN   AND    LEAD  163 

Experiment  238  —  Qualitative  Analysis  of  a  Solution  of 
Lead,  Silver,  and  Mercury  (ous) 

MATERIALS.  —  Solution    containing    lead    nitrate,   silver    nitrate,  and 
mercurous  nitrate. 

Add  dilute  hydrochloric  acid  drop  by  drop  to  the  solution  until 
precipitation  eases.  Allow  the  mixture  of  precipitated  chlorides  to 
settle,  pour  off  the  liquid  carefully  (see  Int.  §  6  (i)(a)),  add  about  15 
cubic  centimeters  of  water,  and  boil.  Filter,  and  test  the  nitrate  for 
lead.  Wash  the  precipitate  with  hot  water  until  the  wash  water  does 
not  give  a  test  for  lead.  Pierce  a  hole  in  the  point  of  the  filter  paper 
with  a  glass  rod,  and  wash  the  mixed  precipitates  of  silver  and  mer- 
curous chlorides  into  a  test  tube  with  dilute  ammonium  hydroxide. 
Warm  gently  and  shake.  Filter,  and  test  the  nitrate  for  silver.  The 
black  residue  is  a  sufficient  test  for  mercury.  Its  presence  may  be 
confirmed  thus:  Dissolve  the  black  precipitate  in  a  very  little  aqua 
regia,  dilute  with  water,  and  add  a  clean  copper  wire;  remove  the 
wire  in  a  few  minutes,  wipe  gently,  and  mercury  will  be  seen  on  the 
wire  as  a  bright  silvery  coating. 


CHROMIUM  —  MANGANESE 

Experiment  239  —  Tests  for  Chromium 

MATERIALS.  —  Borax,  chrome  alum,  potassium  carbonate,  potassium 
nitrate,  acetic  acid,  nitric  acid,  sodium  hydroxide  solution,  lead 
nitrate  solution,  potassium  dichromate  solution,  test  wire,  piece  of 
porcelain,  forceps. 

(a)  Apply  the  borax  bead  test  to  chrome  alum. 

(b)  Mix   equal   small   quantities   of  potassium   carbonate, 
potassium  nitrate,  and  powdered  chrome  alum,  place  the  mix- 
ture on  a  piece  of  porcelain,  hold  it  with  the  test  tube  holder 
in  the  Bunsen  flame,  and  heat  it  until  the  mixture  fuses. 
A  yellow  mass,  due  to  the  presence  of  potassium  chromate, 
results.    If  the  color  is  not  decided,  dissolve  the  mass  in  water, 
add  acetic  acid,  slowly  at  first,  and  boil  until  all  the  carbon 
dioxide  is  expelled.    Add  a  few  drops  of  lead  nitrate  solution 
to  a  portion,  and  yellow  lead  chromate  is  precipitated.     (If 
the  precipitate  is  white,  it  is  lead  carbonate,  and  shows  that 
not  all  the  potassium  carbonate  was  decomposed,  as  intended.) 

(c)  Proceed  as  in  Exp.  233  (e),  using  potassium  chromate  or 
dichromate  solution.      State  the  result. 

Experiment  240  —  Properties  of  Potassium  Chromate 
and  Dichromate 

MATERIALS.  —  Potassium    chromate    and    dichromate,     concentrated 
hydrochloric  acid,  potassium  hydroxide  solution. 

(a)  Make   a   dilute   solution   of  potassium   chromate   and 
dichromate,  and  compare  the  colors.     Save  for  (b)  and  (c). 

(b)  Add  a  few   drops   of   concentrated   hydrochloric    acid 
to.  the  solution  of   potassium  chromate  prepared  in  (a),  and 
observe    the    change  in  color.     Describe  it.      Compare  with 
the  color  of    the  potassium    dichromate    solution.      Draw  a 
conclusion. 


CHROMIUM  —  MANGANESE  1 65 

(c)  Add  potassium  hydroxide  solution  to  the  solution  of 
potassium  dichromate  prepared  in  (a)  until  a  change  of  color 
is  produced.    Describe  the  color.    Compare  with  the  potassium 
chromate  solution.     Draw  a  conclusion. 

(d)  Add  a  few  drops  of  concentrated  hydrochloric  acid  to 
powdered  potassium  chromate  and  dichromate  in  separate 
test  tubes.    What  gas  is  evolved?    By  what  chemical  change 
was  it  produced? 

Experiment  241  —  Tests  for  Manganese 

MATERIALS.  —  Manganese  dioxide,  potassium  carbonate,  potassium 
nitrate,  ammonium  sulphide,  manganese  sulphate  solution,  hydro- 
chloric acid,  acetic  acid,  ammonium  hydroxide. 

(a)  Subject  a  minute  quantity  of  manganese  dioxide  (or 
any  other  manganese  compound)  to  the  borax  bead  test,  and 
note  the  color  of  the  bead  after  heating  in  each  flame. 

(b)  Fuse,  on  a  piece  of  porcelain,  a  little  manganese  dioxide 
mixed  with  potassium  carbonate  and  potassium  nitrate.    The 
green  color  of  the  mass  is  due  to  potassium  manganate. 

(c)  Add  ammonium  sulphide  to  manganese  sulphate  solu- 
tion.    The  flesh-colored  precipitate  is  manganese  sulphide. 
Compare  with  other  sulphides  as  to  color  (see  Exps.  161, 163). 

SUPPLEMENTARY  EXPERIMENTS 

Experiment  242  —  Reduction  of  Chromates  to  Chromic 
Compounds 

MATERIALS.  —  Potassium    dichromate    solution,   concentrated    hydro- 
chloric acid,  alcohol. 

Add  to  a  few  cubic  centimeters  of  potassium  dichromate  solution  a 
little  concentrated  hydrochloric  acid  and  a  few  drops  of  alcohol. 
Warm  gently.  Two  important  changes  occur.  The  chromate  is 
reduced  to  chromic  chloride  which  colors  the  solution  green;  the  alcohol 
is  oxidized  to  aldehyde,  which  is  detected  by  its  peculiar  odor. 


1 66  CHEMISTRY 

Experiment  243  —  Preparation  and  Properties  of  Chromic 
Hydroxide 

MATERIALS.  —  Solutions  of  sodium  hydroxide  and  chrome  alum. 

Add  a  little  sodium  hydroxide  solution  to  a  solution  of  chrome 
alum.  The  precipitate  is  chromic  hydroxide.  Describe  it.  Add  an 
excess  of  sodium  hydroxide  solution,  and  shake.  Describe  the  result. 
Boil,  and  state  the  result.  (Compare  with  Exp.  206.) 

Experiment  244  —  Oxidation  with  Potassium  Permanganate 

MATERIALS.  —  Potassium  permanganate,   sulphuric  acid,   ferrous  sul- 
phate, filter  paper. 

(a)  Add  a  few  drops  of  sulphuric  acid  to  a  weak  solution  of  ferrous 
sulphate;  then  add,  drop  by  drop,  a  dilute  solution  of  potassium 
permanganate.  The  potassium  permanganate  oxidizes  the  ferrous 
to  ferric  sulphate;  its  color  is  changed,  owing  to  the  loss  of  oxygen 
and  transformation  into  other  compounds. 

(6)  Boil  a  piece  of  filter  paper  in  a  dilute  solution  of  potassium 
permanganate.  Describe  and  explain  the  result. 


LABORATORY   EQUIPMENT 


The  lists  given  below  include  the  apparatus,  chemicals,  and  supplies 
needed  for  the  experiments  in  this  book.  Quantities  and  prices  have 
been  omitted  in  justice  to  teachers,  dealers,  and  the  author.  The 
author,  at  his  own  suggestion,  has  lodged  with  the  L.  E.  Knott 
Apparatus  Co.,  15  Harcourt  Street,  Boston,  Mass.,  information 
regarding  the  quantities  needed  for  classes  of  varying  size.  It  is 
hoped  that  teachers  who  desire  information  will  correspond  with 
the  dealer  when  preparing  orders.  The  author  takes  this  opportu- 
nity to  say  he  has  no  financial  connection  with  any  dealer  in  scientific 
supplies. 

List  A  —  Individual  Apparatus 

This  list  includes  the  apparatus  constantly  used,  and  each  student 
should  be  provided  with  a  set. 

Blowpipe. 

Blowpipe  tube. 

Bottles,  wide  mouth,  250  cc. 

Bottle,  generator,  250  cc. 

ft.  Copper  wire,  No.  20. 

Bunsen  burner. 

Deflagrating  spoon. 

Erlenmeyer  flask,  250  cc. 

Evaporating  dish,  2\  in. 

oo  Filter  papers,  4  in. 

Forceps,  iron. 

Funnel,  2\  in. 

4  Glass  plates,  4  x  4  in. 
6  in.  Glass  rod. 

5  ft.  Glass  tubing.1 

i  Graduated  cylinder,  25  cc. 

i  Iron    stand,    clamp    (medium). 


Mortar  and  pestle,  3  in. 
Pinch  clamp  (Mohr's).4 
Pneumatic  trought  complete.2 
Rubber  stopper,  one-hole.3 
Rubber  stopper,  two-hole.3 

2  ft.  Rubber  burner  tubing,  \  in. 

6  in.   Rubber    connecting    tubing, 
T3<r  in. 

6  in.  Rubber     pressure     tubing, 
Ttin.4 

6  Test  tubes,  6  x  \  in. 


3  Test  tubes,  8  x  i  in. 
Test  tube  brush. 
Test  tube  holder. 
Test  tube  rack. 
Thistle  tube. 4 
Wire  gauze,  iron,  4  x  4  in. 
ring  (3  in.). 

1  To  fit  the  rubber  stoppers. 

2  Indurated  fiber  keeler  No.  4  and  shallow  flower  pot,  4  in. 

3  To  fit  the  8xi  in.  test  tube.     Such  stoppers  also  fit  the  average 
250  cc.  Erlenmeyer  flask  and  the  250  cc.  generator  bottle  (if  the  neck 
is  small);   they  also  fit  the  2500  cc.  bottle  (acid  bottle) — see  List  B. 

4  See  Exp.  10.   Straight  stem  of  thistle  tube  must  fit  the  rubber  stopper. 


i68 


CHEMISTRY 


List  B  —  Special  Apparatus 

This  list  includes  apparatus  used  infrequently  or  of  such  a  nature 
that  it  may  be  used  by  a  class  or  groups.  Numbers  in  parentheses 
refer  to  experiments. 

i  Battery,  4  or  more  cells  (30,  70, 
72,  etc.). 

1  Battery  jar,  4-3-  in.  diam.  (70,  72). 

2  Beakers,  No.  3  (63,  78). 
2  Bottles,  2500  cc.  (13). 

2  Burettes,  50  cc.  (78). 

i  Burner,  acetylene  (87). 

i  Chlorine  tube  (29). 

i  Clamp,  medium  (extra)  (78). 

i  Condenser,  complete  (26,  127). 

i  Crucible   and   cover,   porcelain, 

No.  o  (9,  32,  81). 
i  Crucible,  Hessian  (sand),  4  in. 

deep  (80,  208). 
i  Crucible,  iron,  60  cc.  (55). 
i  Crucible  block,  wood,  4  x  4  x  i  in. 

(9,  32). 
i  Dish,  lead  (147,  150,  152). 

1  Electric  bell  (70,  72). 

2  Electrodes,  carbon  (70,  72). 

i  Electrolysis  apparatus  (30,  75). 


i  Flask,  500  cc.  (26,  127). 

i  Hydrometer  (135). 

i  Jar,  12x4  in.  (62,  65). 

i  Magnet  (216). 

6  in.  Nichrome  wire  (Int.  6  (4)). 

i  Pan,  iron,  4  in.  (204). 

i  Sq.  in.  Platinum  foil  (30). 

3  in.  Platinum  wire  (Int.  5  (4)). 

i  Retort,  250  cc.  (47). 

i  Screw,  Hofmann  (13). 

i  Thermometer,  — 10    to    110°    C. 

(frequently), 
i  Triangle  (9,  32). 
i  Tripod  (9,  32). 
i  Tube    graduated,    100    cc.    (62, 

65). 

i  U-tube,  4  in.  (71,  77). 
i  Welsbach    burner    and    mantle 

(102). 
i  Wing-top  burner  (Int.  3  (6)). 


List  C  —  General  Apparatus 
Numbers  in  parentheses  refer  to  experiments. 


Balance  (61-65). 

Barometer  (13,  61,  62,  64,  65). 

set  Cork  borers  (Int.  6  (3)). 

File,  round  ^(Int.  5,  (3)). 

File,  triangular  (Int.  3  (a)). 
2  Cylinders,    graduated,    250    cc. 

and  500  cc. 
i  Microscope  (108). 


i  Scales. 

i  set  Weights  for  balance,   i  mg. 

to  50  gm. 
i  set  Weights  for  scales,  i  gm.  to 

500  (or  1000)  gm. 
Wooden  blocks,  6  x  6  x  i  and  4  x 

4x5  (with  %  in.  hole  in  center) 

(33,  35)- 


LABORATORY   EQUIPMENT 


169 


List  D  —  Chemicals 

This  list  includes  the  chemicals  (except  those  in  List  E)  needed 
for  the  experiments  in  this  book.  Numbers  in  parentheses  refer  to 
experiments  in  which  the  chemicals  are  used  infrequently  or  in  small 
quantities. 


Acid,  acetic. 

hydrochloric. 

metaphosphoric  (160). 

nitric. 

orthophosphoric  (159). 

oxalic  (91). 

pyrogallic  (57). 

sulphuric. 
Alcohol,  ethyl. 

methyl  (109,  128). 
Alum,  chrome  (239,  243). 

potassium. 
Albumin. 
Aluminium,  bronze  (188). 

granulated  (208). 

lump  (205). 

sheet. 

wire  (13). 

sulphate. 

Ammoniacal  liquor  (53,  99). 
Ammonium  carbonate   (161,   177, 
1 80). 

chloride. 

dichromate  (52,  76). 

hydroxide. 

molybdate  (117,  119,  159). 

nitrate  (56). 

oxalate. 

polysulphide  (161,  237). 

sulphate  (53). 

sulphide  (241). 
Aniline  (18). 
Antimony  (165,  166). 

chloride  (162). 

See  Tartar  emetic. 
Arsenic  trioxide  (139,  161). 
Asbestos  (13). 


Baking  powder  (List  E). 

alum  (113,  215). 
Barium  chloride. 

dioxide  (8,  31,  208). 

hydroxide  (60). 

nitrate  (52). 

peroxide  (see  dioxide). 
Bauxite,  red  (220). 
Bismuth    (167,  168). 
Bleaching  powder  (34,  197). 
Bone  ash  (159,  197). 
Borax. 
Brass  (188). 

Cadmium  chloride  (231). 
Calcium  (12,  16,  29,  65,  195). 

carbide  (87). 

carbonate  (marble). 

chloride. 

fluoride  (150,  152). 

oxide  (lime). 

sulphate. 
Camphor  (109). 
Carbon  disulphide. 

tetrachloride  (114). 
Carborundum,  lump  (93). 

powder  (150). 
Casein  (103). 
Celluloid  (126). 
Cement  (197). 
Chalk  (86). 
Charcoal,  animal. 

lump. 

powder. 

Chloroform  (66). 
Chrome  alum.     See  Alum. 

yellow  (236). 


i  yo 


CHEMISTRY 


Clay  (207). 
Cobalt  chloride. 

nitrate  (76). 
Cochineal  (210). 
Collodion  (126). 
Copper,  borings. 

wire  (No.  20). 

bromide  (76). 

nitrate  (51,  76). 

sulphate. 
Cream  of  tartar  (44). 

Dextrin  (125). 
Ether  (114,  129). 

Fehling's  solution  (  see  Exp.  106). 
Ferric  ammonium  citrate  (221). 
Ferric  and  ferrous  salts  (see  Iron) 
Formalin  (128). 
Fusible  metal  (2,  169). 

Gasoline  (114). 

Gelatin. 

German  silver  (188). 

Glass  wool  (13). 

Glycerin  (18). 

Glucose.     See  Grape  sugar. 

Gold  leaf. 

Grape  sugar. 

Graphite,  artificial  (88). 

native  (88). 
Guncotton  (126). 

Hydrogen  dioxide  (8). 

Iodine. 

Infusorial  earth  (150). 

Iron  filings  (54). 

powder. 

thread  (Steel  wool). 

chloride  (ic). 

sulphate  (ous). 

sulphide  (ous). 


Lead,  sheet. 

tea. 

acetate. 

dioxide  (8,  41). 

nitrate. 

monoxide  (litharge). 

red  (tetroxide),  (41,  236). 

white  (carbonate),  (86,  236). 
Lignite  (81). 

Lime.     See  Calcium  oxide. 
Litmus  paper. 

solution  (71). 

Magnesium,  powder  (32). 

ribbon. 

carbonate  (89). 

chloride. 

oxide  (224). 

sulphate. 
Manganese  dioxide. 

sulphate  (241). 
Mercury  (229). 

chloride  (ic). 

nitrate  (ic),  (226). 

nitrate  (ous),  (226,  238). 

oxide  (8). 

Nickel  chloride  (76,  222). 
sulphate  (76). 

Paraffin  wax  (152). 
Phenol-phthalein  solution  (78,  87, 

90). 
Phosphorus,  red  (164). 

yellow  (164). 
Plaster  of  Paris  (97). 
Potash  (178). 
Potassium  (16,  173). 

bromide. 

carbonate. 

chlorate  (cryst.). 

chlorate  (powd.). 

chloride. 


LABORATORY   EQUIPMENT 


171 


Potassium  (16,  173). 

chromate. 

cyanide  (186). 

dichromate. 

ferricyanide. 

ferrocyanide. 

hydroxide. 

iodide. 

nitrate. 

nitrite  {223). 

perchlorate  (66). 

permanganate. 

sulphate. 

sulphocyanate  (218,  220). 
Pumice  (150). 

Rochelle  salt  (see  Exp.  106). 
Rosin  (109). 

Shellac  (powd.),  (204). 
Silicon  (146). 
Silver  nitrate. 

sulphate  (73). 
Soda-lime. 
Sodium,  metal. 

acetate  (69). 

ammonium  phosphate  (160). 

bicarbonate. 

carbonate. 

chloride. 

cobaltinitrite  (174). 

dichromate  (76). 

hydroxide. 

nitrate. 

nitrite  (45,  69). 

peroxide  (8). 


Sodium,  phosphate  (acid),  (69). 

phosphate     (disodium),     (69, 
224). 

silicate  (148). 

sulphate. 

sulphate  (acid),  (69). 

sulphite  (134)- 

thiosulphate    (hyposulphite), 

(28,  193). 
Solder  (236,  237). 
Steel  wool.     See  Iron  thread. 
Stannous  chloride  (232). 
Strontium  nitrate  (200,  204). 
Sulphur,  flowers  (130). 

roll. 

Tannin  (125). 
Tartar  emetic  (139). 
Thermit  (208). 
Tin  foil  (236). 

granulated. 

rod  (234,  235). 
Turpentine  (33). 
Type  metal  (236). 

Whiting  (86,  197). 
Wood's  metal  (169),  (see  Fusible 
metal). 

Zinc  dust  (9). 

granulated. 

sheet. 
Zinc  chloride  (24). 

oxide  (225). 

sulphate. 


172 


CHEMISTRY 


List  E  —  Miscellaneous  Supplies 
Numbers  in  parentheses  refer  to  experiments. 


Apple  (14). 

Baking  powder  (86,  113,  215). 

Bluing  (220). 

Bread  (14,  123,  124). 

Brick. 

Bullets  (236). 

Butter  (114)'. 

Calico  (33). 

Candle,  paraffin  (80). 

Candy  (107). 

Celery  (14). 

Cloth,  colored  (34). 

unbleached  (34). 
Coal,  hard  (81). 

soft  (8 1,  94). 
Coke  (99). 

Corks,  assorted  (see  Exp.  82). 
Cotton. 

Cotton-seed  oil  (115). 
Cracker  (14,  124). 
Cranberry  (14). 
Emery  paper  (Int.  3). 
Excelsior  (25). 
Fertilizer  (159). 
Flour  (120). 
Flower  pot  (220). 
Gas  carbon  (99). 
Gelatin  (103). 
Grape  juice  (44). 
Hay  (25). 
Iron,  cast  (216). 

nails. 

rust  (220). 

steel  (216). 

wrought  (216). 
Jam  (107). 


Jelly  (44,  107). 

Joss  sticks  (frequently  used). 

Kerosene  (18). 

Lard  (114,  115). 

Leather  (103). 

Lemon  juice  (44). 

Maple  sugar  (107). 

Matches. 

Meat  (14,  103). 

Milk  (44). 

Molasses  (107). 

Mortar,  old  (86,  197). 

Mustard  (103). 

Pickles  (44). 

Plaster  (197). 

Potato  (14,  108,  123). 

Quill  toothpick  (48). 

Raisins  (25). 

Rice  (123). 

Rock,  rusty  (220). 

Sand. 

Shot  (236). 

Sirup,  table  (107). 

Soap. 

Starch  (108). 

Sugar. 

Tallow  (114). 

Tar  (99). 

Taper,  wax  (10). 

Thread. 

Tin,  sheet  (220). 

Tooth  powder. 

Vinegar  (112). 

Wood  ashes  (44). 

Yeast  (127). 


LABORATORY   EQUIPMENT  173 

List  F  —  Emergency  Supplies  (See  Int.  6) 

Absorbent  cotton,  i  oz.  Medicine  dropper. 

Bandages '(2  in.).  Mustard,  i  oz. 

Blanket.  Pins,  i  paper. 

Boric  acid  solution,  i  oz.  Sand  and  scoop. 

Camphor  solution,  i  oz.  Scissors. 

Collodion,  \  oz.  Smelling  salts,  i  bottle. 

Carron  oil,  i  oz.  Sodium  bicarbonate,  i  oz. 

Court  plaster,  i  book.  Thread,  i  spool. 

Fire  extinguisher.  Vaseline,  i  oz. 

List  G  —  Solutions 

The  solutions  needed  for  most  of  the  experiments  in  this  book  are 
approximately  10  per  cent  (unless  otherwise  specified)  except  the 
following: 

Ammonium  hydroxide,  i  vol.  to  3  vols.  water. 
Ammonium  molybdate.     Dissolve  15  gm.  of  the  salt  in  100  cc.  water 

and  pour  this  solution  into  100  cc.  nitric  acid  (i  vol.  acid  to  i  vol. 

water). 

Ammonium  oxalate,  5  per  cent. 
Ammonium    polysulphide.     Add    sulphur     to     ammonium    sulphide 

solution. 

Ammonium  sulphide,  i  vol.  to  i  vol.  water. 
Barium  chloride,  5  per  cent.     Use  distilled  water. 
Barium  hydroxide,  i  per  cent.     Use  clear  solution. 
Battery  solution  (Grenet).    Dissolve  103  gm.  powd.  potassium  dichro- 

mate  in  1000  cc.  water  and  slowly  add  103  gm.  cone,  sulphuric  acid 

with  constant  stirring. 

Chlorine  water.     Saturate  water  with  the  gas. 
Cobaltous  nitrate,  5  per  cent. 
Cochineal.     Grind  a  little  cochineal  with  water,  dilute  as  desired,  and 

filter. 

Ferric  chloride,  5  per  cent. 

Ferrous  sulphate,  5  per  cent.     Must  be  freshly  prepared. 
Hydrochloric  acid,  dilute,  i  vol.  to  5  vols.  water. 
Lead  salts.    Use  distilled  water  or  filter. 
Limewater.    Slake  lime,  add   considerable  water,  shake   occasionally, 

siphon  off  clear  liquid. 
Litmus.     As  under  Cochineal. 
Mercuric  chloride,  5  per  cent.     POISON. 
Mercurous  nitrate,  5  per  cent.     Add  a  little  mercury. 


174  CHEMISTRY 

Nitric  acid,  dilute,  i  vol.  to  5  vols.  water. 

Phenol-phthalein.     Dissolve  i  gm.  in  100  cc.  alcohol  and  dilute  slightly 

with  water. 

Potassium  permanganate,  5  per  cent. 
Potassium  sulphocyanate,  5  per  cent. 
Silver  nitrate,  5  per  cent.     Use  distilled  water. 
Silver  sulphate,  .5  per  cent.     Use  distilled  water. 
Sodium  cobaltinitrite.     Dissolve  10  gm.  sodium  nitrite  in  20  cc.  water; 

dissolve  6  cc.  acetic  acid  and  i  gm.  cobaltous  nitrate  in  15  cc.  water. 

Mix  these  solutions  and  filter. 
Stannous  chloride.     Reduce  cone,   hydrochloric  acid  with    tin,  dilute 

with  water,  and  keep  tin  in  this  solution. 
Sulphuric  acid,  dilute.     Slowly  pour  i  vol.  acid  into  5  vols.  of  water, 

stirring  constantly. 


UNIVERSITY  OF  CALIFORNIA  LIBRARY 
BERKELEY 

Return  to  desk  from  which  borrowed. 
This  book  is  DUE  on  the  last  date  stamped  below. 


24Apr'59DF 
REC'D  LD 

APR  10  1959 


WT.UHW 


LD  21-100m-ll,'49(B7146sl6)476 


YB  16942 

TABLE  OF  ATOMIC  WEIGHTS 


Element 

1 

a 

i 

Inter- 
national 

|| 

<j 

Element 

1 

I 

Jl'il 

a 

Approxi- 
mate 

Aluminium 

Al 

27.1 

27 

Molybdenum 

Mo 

96.0 



Antimony 

Sb 

120.2 

120 

Neodymium 

Nd 

144.3 

— 

Argon 

A 

39.88 

40 

Neon 

Ne 

20.2 

20 

Arsenic 

As 

74.96 

75 

Nickel 

Ni 

58.68 

58.5 

Barium 

Ba 

137.37 

137 

Niton 

Nt 

222.4 

— 

Bismuth 

Bi 

208.0 

208 

Nitrogen 

N 

14.01 

14 

Boron 

B 

11.0 

11 

Osmium 

Os 

190.9 

— 

Bromine 

Br 

79.92 

80 

Oxygen 

O 

16.00 

16 

Cadmium 

Cd 

112.40 

112 

Palladium 

Pd 

106.7 

— 

Caesium 

Cs 

132.81 

— 

Phosphorus 

P    . 

31.04 

31 

Calcium 

Ca 

40.07 

40 

Platinum 

Pt 

195.2 

195 

Carbon 

C 

12.00 

12 

Potassium 

K 

39.10 

39 

Cerium 

Ce 

140.25 

— 

Praseodymium 

Pr 

140.6 

— 

Chlorine 

Cl 

35.46 

35.5 

Radium 

Ra 

226.4 

226.5 

Chromium 

Cr 

52.0 

52 

Rhodium 

Rh 

102.9 

— 

Cobalt 

Co 

58.97 

59 

Rubidium 

Rb 

85.45 

— 

Columbium 

Cb 

93.5 

— 

Ruthenium 

Ru 

101.7 

— 

Copper 

Cu 

63.57 

63.5 

Samarium 

Sa 

150.4 

— 

Dysprosium 

Dy 

162.5 

— 

Scandium 

Sc 

44.1 

— 

Erbium 

Er 

167.7 

— 

Selenium 

Se 

79.2 

— 

Europium 

Eu 

152.0 

— 

Silicon 

Si 

28.3 

28 

Fluorine 

F 

19.0 

19 

Silver 

Ag 

107.88 

108 

Gadolinium 

Gd 

157.3 

— 

Sodium 

Na 

23.00 

23 

Gallium 

Ga 

69.9 

— 

Strontium 

Sr 

87.63 

87.5 

Germanium 

Ge 

72.5 

— 

Sulphur 

S 

32.07 

32 

Glucinum 

Gl 

9.1 

— 

Tantalum 

Ta 

181.5 



Gold 

Au 

197.2 

197 

Tellurium 

Te 

127.5 



Helium 

He 

3.99 

4 

Terbium 

Tb 

159.2 



Holmium 

Ho 

163.5 

— 

Thallium 

Tl 

204.0 



Hydrogen 

H 

1.008 

1 

Thorium 

Th 

232.4 



Indium 

In 

114.8 

— 

Thulium 

Tm 

168.5 



Iodine 

I 

126.92 

127 

Tin 

Sn 

119.0 

119 

Iridium 

IT 

193.1 

— 

Titanium 

Ti 

48.1 

Iron 

Fe 

55.84 

66 

Tungsten 

W 

184.0 



Krypton 

Kr 

82.92 

— 

Uranium 

U 

238.5 



Lanthanum 

La 

139.0 

— 

Vanadium 

V 

51.0 



Lead 

Pb 

207.10 

207 

Xenon 

Xe 

130.2 



Lithium 

Li 

6.94 



Ytterbium 

Yb 

172.0 



Lutecium 

Lu 

174.0 

— 

Yttrium 

Yt 

89.0 



Magnesium 

Mg 

24.32 

24 

Zinc 

Zn 

65.37 

65 

Manganese 

Mn 

54.93 

55 

Zirconium 

Zr 

90.6 

Mercury 

Hg 

200.6 

200.5 

